MEMCAL 


Gift  of 
FRANK  EENMAN,   M.D 


AN    INTRODUCTORY    COURSE 


OF 


QUANTITATIVE 
CHEMICAL   ANALYSIS, 

WITH 

EXPLANATORY     NOTES 

AND 

STOICHIOMETR1CAL     PROBLEMS. 


BY 


HENRY    P.    TALBOT,    PH.D., 

PROFESSOR    OF    ANALYTICAL    CHEMISTRY     IN     THE    MASSACHUSETTS 
INSTITUTE    OF    TECHNOLOGY. 


THIRD  EDITION. 
REVISED  AND  ENLARGED. 


NEW    YORK: 
THE    MACMILLAN    CO. 

LONDON  :  MACMILLAN  &  Co.,  Ltd. 
1899. 


COPYRIGHT,  1897, 
BY  HENRY  P.  TAL'BOT. 


PREFACE. 


THIS  Introductory  Course  of  Quantitative  Analysis  has  been 
prepared  to  meet  the  needs  of  students  who  are  just  entering 
upon  the  subject,  after  a  course  of  qualitative  analysis.  It  is 
primarily  intended  to  enable  the  student  to  work  successfully 
and  intelligently,  without  the  necessity  for  a  larger  measure 
of  personal  assistance  and  supervision  than  can  reasonably  be 
given  to  each  member  of  a  large  class.  To  this  end  the  direc- 
tions are  given  in  such  detail  that  there  is  very  little  oppor- 
tunity for  the  student  to  go  astray ;  but  the  manual  is  not,  the 
author  believes,  on  this  account  less  adapted  for  use  with  small 
classes,  where  the  instructor,  by  greater  personal  influence,  cart 
stimulate  independent  thought  on  the  part  of  the  pupil. 

The  method  of  presentation  of  the  subject  is  that  suggested 
by  Prof.  A.  A.  Noyes'  excellent  manual  of  Qualitative  Analysis. 
For  each  analysis  the  procedure  is  given  in  considerable  detail, 
and  this  is  accompanied  by  explanatory  notes,  which  are  believed 
to  be  sufficiently  expanded  to  enable  the  student  to  understand 
fully  the  underlying  reason  for  each  step  prescribed.  The  use 
of  the  book  should  nevertheless  be  supplemented  by  classroom 
instruction,  mainly  of  the  character  of  recitations,  and  the  student 
should  be  taught  to  consult  the  larger  works,  such  as  those  of 
Fresenius,  Mohr,  and  Sutton.  The  general  directions  of  Part  I 
are  intended  to  emphasize  those  matters  upon  which  the  begin- 
ner in  quantitative  analysis  must  bestow  special  care,  and  to 
offer  some  helpful  suggestions.  The  student  can  hardly  be  ex- 
pected to  appreciate  the  force  of  all  the  statements  contained  in 
these  directions,  or,  indeed,  to  retain  them  all  in  the  memory- 
after  a  single  reading,  but  the  instructor,  by  frequent  reference 


4  PREFACE. 

to  special  paragraphs  as  suitable  occasion  presents  itself,  can 
soon  render  them  familiar  to  the  student. 

The  analyses  selected  for  practice  are  those  comprised  in  the 
course  of  "  preliminary  quantitative  analysis  "  at  the  Massachu- 
setts Institute  of  Technology,  and  have  been  chosen,  after  an 
experience  of  some  years,  as  affording  the  best  preparation  for 
.more  advanced  work,  and  as  satisfactory  types  of  gravimetric 
and  volumetric  methods.  From  the  latter  point  of  view,  they 
also  seem  to  furnish  the  best  insight  into  quantitative  analysis 
for  those  students  who  can  devote  but  a  limited  time  to  the  sub- 
ject, and  who  may  never  extend  their  study  beyond  the  field  cov- 
ered by  this  manual.  The  author  has  had  opportunity  to  test  the 
efficiency  of  the  course  for  use  with  such  students,  and  has 
found  the  results  satisfactory. 

In  place  of  the  usual  custom  of  selecting  simple  salts  as  mate- 
rial for  preliminary  practice,  it  has  been  found  advantageous  to 
substitute,  in  most  instances,  approximately  pure  samples  of  ap- 
propriate minerals  or  industrial  products.  The  difficulties  are 
not  greatly  enhanced,  while  the  student  gains  in  practical 
'experience. 

It  has  been  found  expedient  with  large  classes,  to  allow  the 
whole  class  to  work  simultaneously  upon  the  same  procedure,  for 
example,  that  for  the  determination  of  chlorine  in  sodium  chlo- 
ride, since  classroom  instruction  can  then  be  made  more  effective. 
Each  individual  is,  however,  permitted  to  work  as  rapidly  as  his 
capacity  admits,  and  to  such  students  as  exhibit  unusual  facility  in 
manipulation,  extra  analyses  are  assigned. 

The  author  has  been  unable  to  find  any  work  in  which  such 
:stoichiometrical  problems  as  are  constantly  met  with  in  the  experi- 
ence of  an  analyst  are  dealt  with  in  such  detail  as  to  enable  the 
student  to  fully  understand  the  underlying  principles.  A  chap- 
ter has  therefore  been  added  in  which  such  problems  are  pre- 
sented, and  the  solutions  of  certain  typical  cases  are  explained. 
A  table  of  atomic  weights  and  a  table  of  four-place  logarithms  are 
appended  for  convenience. 


PREFACE.  5 

The  analytical  procedures,  as  detailed  in  this  manual,  are  those 
dictated  by  experience  in  the  laboratories  of  the  Institute  as  giv- 
ing the  best  results  for  purposes  of  instruction.  The  author  has 
received  many  suggestions  from  standard  works,  but  no  attempt 
has  been  made  to  enumerate  these  in  full ;  nor,  on  the  other  hand, 
is  any  considerable  credit  claimed  for  originality.  Criticisms  or 
corrections  will  be  welcomed. 

A  knowledge  of  the  modern  theories  of  solutions  has  become 
so  important  a  possession  for  even  the  beginner  in  quantitative 
analysis,  if  he  is  to  view  his  work  from  a  scientific  standpoint, 
that  it  has  seemed  desirable  to  add  a  chapter  to  the  present 
edition,  in  which  these  theories  are  briefly  stated.  No  attempt 
has  been  made  to  treat  the  subject  exhaustively,  and  the  instructor 
will  doubtless  find  it  necessary  and  desirable  to  present  much  in 
the  way  of  additional  detail  and  illustration.  The  student  should 
also  be  urged  to  think  and  read  for  himself. 

A  few  determinations  have  been  added  to  the  list  of  gravimetric 
analyses,  and  the  text  has  been  generally  revised. 

The  author  wishes  to  express  his  appreciation  of  the  kindly 
reception  accorded  the  earlier  editions  of  this  manual,  and  his 
obligations  to  Prof.  A.  A.  Noyes,  and  to  Drs.  W.  H.  Walker  and 
F.  J.  Moore,  Instructors  in  Analytical  Chemistry  at  the  Insti- 
tute, for  helpful  suggestions  and  much  kind  assistance  in  its 
preparation. 


HENRY  P.  TALBOT. 


MASSACHUSETTS  INSTITUTE  OF  TECHNOLOGY, 
January,  1899. 


PART   I. 


INTRODUCTION. 


A  COMPLETE  chemical  analysis  of  a  body  of  unknown 
composition  involves  the  recognition  of  its  component  parts 
by  the  methods  of  qualitative  analysis,  and  the  determina- 
tion of  the  proportions  in  which  these  components  are  pres- 
ent by  the  processes  of  quantitative  analysis.  The  quali- 
tative examination  is  indispensable  if  intelligent  and  proper 
provisions  are  to  be  made  for  the  separation  of  the  various 
constituents,  under  such  conditions  as  shall  insure  accurate, 
quantitative  estimations. 

It  is  assumed  that  the  operations  of  qualitative  analysis 
are  familiar  to  the  student,  who  will  find  that  the  reactions 
made  use  of  in  quantitative  processes  are  not  infrequently 
those  employed  for  the  qualitative  detection  of  the  same 
element ;  but  it  should  be  noted  that  the  conditions  must 
now  be  regulated  with  greater  care,  and  in  such  a  manner 
as  to  insure  the  most  complete  separations  possible.  For 
example,  in  the  qualitative  detection  of  sulphates  by  pre- 
cipitation as  barium  sulphate  from  acid  solution,  it  is  not 
necessary,  in  most  instances,  to  regard  the  solubility  of  the 
sulphate  in  hydrochloric  acid,  while  in  the  quantitative  de- 
termination of  sulphates  by  this  reaction,  this  solubility  be- 
comes an  important  consideration.  The  operations  of  qual- 
itative analysis  are,  therefore,  the  more  accurate,  the  nearer 
they  are  made  to  conform  to  quantitative  conditions. 

The  methods  of  quantitative  analysis  are  sub-divided,  ac- 
cording to  their  nature,  into  those  of  gravimetric  and  vol- 
umetric analysis.  In  gravimetric  processes  the  constituent 
to  be  determined  is  isolated  in  the  form  of  some  compound 
possessing  a  well-established  and  definite  composition,  which 
can  be  readily  and  completely  separated  by  filtration,  and 


8  GENERAL    DIRECTIONS. 

weighed  either  directly,  or  after  .ignition.  From  the  weight 
of  this  body  the  amount  of  the  constituent  in  question  is 
determined. 

In  volumetric  analysis,  instead  of  the  final  weighing  of  a 
definite  body,  a  well-defined  reaction  is  caused  to  take  place, 
wherein  the  reagent  is  added  in  the  form  of  a  solution,  of 
which  the  strength  (and  hence  the  value  for  the  reaction  in 
question)  is  accurately  known.  The  volume  of  this  solution 
required  to  complete  the  reaction  then  becomes  a  measure 
of  the  substance  acted  upon.  An  example  will  make  the 
distinction  clear.  The  percentage  of  chlorine  in  a  sample  of 
sodium  chloride  may  be  determined  by  precipitation  of  the 
chlorine  from  a  weighed  portion  as  silver  chloride,  which 
is  separated  by  filtration,  ignited,  and  weighed  (a  gravime- 
tric process)  ;  or  the  sodium  chloride  may  be  dissolved  in 
water,  and  a  solution  of  silver  nitrate,  containing  an  accu- 
rately known  amount  of  the  silver  salt  in  each  cubic  centi- 
meter, may  be  cautiously  added  until  precipitation  is  com- 
plete, when  the  amount  of  chlorine  may  be  calculated  from 
the  number  of  cubic  centimeters  of  the  silver  nitrate  solu- 
tion involved  in  the  reaction.  This  is  a  volumetric  process 
and  is  equivalent  to  weighing  without  the  use  of  a  balance. 

Volumetric  methods  are  generally  more  rapid,  and  fre- 
quently capable  of  greater  accuracy,  than  gravimetric  methods. 


GENERAL    DIRECTIONS    FOR    QUANTITATIVE    WORK. 

The  following  suggestions  should  be  carefully  and  thought- 
fully read ;  their  adoption  will  lead  to  work  of  a  high  grade 
of  excellence,  while  their  rejection  may  often  lead  to  unsat- 
isfactory or  careless  work. 

NEATNESS. 

The  laboratory  desk,  and  all  apparatus,  should  be  scrupu- 
lously neat  and  clean  at  all  times.  A  sponge  should  always 
be  ready  at  hand,  and  desk  and  filter-stands  should  be  dry 
and  in  good  order.  Funnels  should  never  be  allowed  to 
drip  upon  the  base  of  the  stand.  Glassware  should  always 
be  wiped  with  a  clean,  lintless  towel  just  before  use. 


WA SH-B  O  TTLES  A ND  DESICCA  TORS,  9 

WASH-BOTTLES. 

Wash-bottles,  for  distilled  water,  should  be  made  from 
flasks  of  about  750  cc.  capacity  and  be  provided  with  grace- 
fully bent  tubes,  which  should  not  be  too  long.  The  jet 
should  be  connected  with  the  tube  entering  the  wash-bottle 
by  a  short  piece  of  rubber  tubing,  in  such  a  way  as  to  be 
flexible,  and  should  deliver  a  stream  about  one  millimeter 
in  diameter.  The  neck  of  the  flask  may  be  wound  with 
twine,  or  covered  with  wash  leather  for  greater  comfort 
when  hot  water  is  used.  It  is  well  to  provide  several 
small  wash-bottles  for  liquids  other  than  distilled  water, 
which  should  invariably  be  clearly  labelled. 

DESICCATORS. 

Desiccators  should  be  filled  with  fused,  anhydrous  calcium 
chloride,  over  which  is  placed  an  iron  triangle  wound  with 
platinum  foil  at  those  points  which  come  into  contact  with 
a  hot  crucible.  The  cover  of  the  desiccator  should  be  made 
air-tight  by  the  use  of  a  thin  coating  of  tallow. 

Pumice  moistened  with  sulphuric  acid  may  be  used  in 
place  of  the  calcium  chloride,  and  is  essential  in  special 
cases,  but  for  most  purposes  the  calcium  chloride,  if  re- 
newed occasionally  and  not  allowed  to  cake  together,  is 
equally  efficient. 

Desiccators  should  never  remain  uncovered  for  any  length 
of  time.  The  dehydrating  agents  rapidly  lose  their  efficiency 
on  exposure  to  the  air. 

CRUCIBLES. 

Platinum  crucibles  should  be  employed  for  all  ignitions 
and  fusions,  when  possible.  All  crucibles,  whether  of  plat- 
inum or  porcelain,  must  be  heated,  and  cooled  in  a  desic- 
cator before  use.  This  is  to  insure  parallel  conditions  in 
separate  weighings,  which  could  not  be  obtained  if  the  cru- 
cible were  cooled  in  contact  with  the  air,  since  a  layer  of 
moisture  is  then  condensed  on  its  surface,  the  amount  vary- 
ing with  the  humidity  of  the  atmosphere.  In  the  dry  air 
of  the  desiccator  this  difficulty  is  avoided. 


io  GENERAL    DIRECTIONS. 

Crucibles  should  be  cleaned,  heated,  and  weighed  before 
each  analysis. 

Platinum  crucibles  should  be  frequently  scoured,  either 
with  sea  sand  or  some  preparation  of  the  general  character 
of  "sapolio."  Constant  heating  causes  a  slight  crystalliza- 
tion of  the  surface  of  the  platinum,  which,  if  not  removed, 
penetrates  into  the  crucible.  Gentle  abrasion  of  the  sur- 
face destroys  the  crystalline  structure  and  prevents  further 
damage.  If  sea  sand  is  used  great  care  is  necessary  to 
keep  it  from  the  desk,  since  beakers  are  easily  scratched 
by  it,  and  subsequently  broken  on  heating. 

Platinum  crucibles  stained  by  iron  may  often  be  cleaned 
by  the  use  of  potassium  acid  sulphate,  or  by  heating  with 
ammonium  chloride.  If  the  former  is  used,  care  should  be 
taken  not  to  heat  so  strongly  as  to  expel  all  of  the  sul- 
phuric acid,  since  the  normal  sulphate  expands  so  rapidly 
on  cooling  as  sometimes  to  burst  the  crucible. 

Bodies  containing  metals  which  might  be  reduced,  with 
the  formation  of  metallic  buttons,  must  not  be  treated  in 
platinum  crucibles.  Fusible  alloys  of  platinum  may  be 
formed  which  ruin  the  crucible.  Compounds  of  phospho- 
rus or  arsenic  must  not  be  heated  under  reducing  condi- 
tions, since  these  elements,  by  contact  with  the  platinum, 
render  it  brittle. 

Liquids  containing  free  chlorine,  aqua  regia,  or  ferric  chlo- 
ride all  exert  a  solvent  action  upon  platinum,  the  ferric  chlo- 
ride to  a  lesser  degree  than  the  others.  Care  must  be  taken 
to  prevent  the  introduction  of  platinum  into  analyses  by  a 
disregard  of  these  facts. 

Caustic  alkalies  and  peroxides  of  the  alkalies  attack  plat- 
inum freely.  Fusions  with  these  fluxes  should  be  made  in 
silver  crucibles. 

EVAPORATION    OF    LIQUIDS. 

Too  great  care  cannot  be  taken  to  prevent  loss  of  solu- 
tions during  processes  of  evaporation,  either  from  too  vio- 
lent ebullition,  from  evaporation  to  dryness  and  spattering, 
or  from  the  evolution  of  gas  during  the  heating.  It  may  be 


REAGENTS.  IX 

stated  in  general  that  evaporation  upon  the  steam  bath  is 
to  be  preferred  to  other  methods  on  account  of  the  im- 
possibility of  loss  by  spattering.  If  the  steam  baths  are 
well  protected  from  dust,  solutions  should  be  left  without 
covers  during  evaporation,  but  solutions  which  are  boiled 
upon  the  hot  plate,  or  from  which  gases  are  escaping,  should 
invariably  be  covered.  In  any  case  a  watch-glass  may  be 
supported  above  the  vessel  by  means  of  a  glass  triangle,  or 
other  similar  device,  and  the  danger  of  loss  of  material  or 
contamination  by  dust  be  thus  avoided. 

It  is  obvious  that  evaporation  is  promoted  by  the  use  of 
vessels  which  admit  of  the  exposure  of  a  broad  surface  to 
the  air. 

Liquids  which  contain  suspended  matter  (precipitates) 
should  always  be  cautiously  treated,  since  the  presence  of 
the  solid  matter  is  frequently  the  occasion  of  violent  "  bump- 
ing," with  consequent  risk  to  apparatus  and  analysis. 

Liquids  should  never  be  transferred  from  one  vessel  to 
another,  nor  to  a  filter,  without  the  aid  of  a  stirring  rod 
held  firmly  against  the  side  or  lip  of  the  vessel.  When  the 
vessel  is  provided  with  a  lip  it  is  not  usually  necessary  to 
use  tallow  or  vaseline  to  prevent  the  loss  of  liquid  by  run- 
ning down  the  side ;  whenever  this  seems  imminent  a  very 
thin  layer  of  tallow,  applied  with  the  finger  to  the  edge  of 
the  vessel,  will  suffice.  The  stirring  rod,  down  which  the 
liquid  runs,  should  never  be  drawn  upward  in  such  a  way 
as  to  allow  the  solution  to  collect  on  the  under  side  of  the 
rim  of  a  beaker. 

REAGENTS. 

All  reagents  should  be  measured,  and  a  record  of  the 
amounts  used  should  be  made  in  the  notebook. 

Whenever  it  is  practicable,  the  amount  of  the  reagent 
required  should  be  calculated,  and  a  large  excess  avoided. 
Many  analyses  are  spoiled  by  a  neglect  of  this  precaution. 

Reagents  should  be  carefully  examined  for  impurities.  If 
these  are  found,  blank  analyses  must  be  made,  using  only 
the  reagents,  and  the  amounts  thus  found  deducted  from 
the  weights  of  contaminated  precipitates.  Under  these  cir- 


I2  GENERAL    DIRECTIONS. 

cumstances,  the  value  of  the  first  suggestion  in  this  para- 
graph is  obvious. 

The  stoppers  of  reagent  bottles  should  never  be  laid  upon 
the  desk,  unless  upon  a  clean  watch-glass  or  paper.  The 
neck  and  mouth  of  all  such  bottles  should  be  kept  scru- 
pulously clean,  and  care  taken  that  no  confusion  of  stoppers 
occurs. 

PRECIPITATION. 

From  theoretical  considerations  it  appears  that  no  sub- 
stance is  to  be  regarded  as  absolutely  insoluble  in  a  spe- 
cific medium,  although  the  solubility  of  many,  which  we 
term  insoluble  bodies,  is  less  than  can  be  measured  by  the 
means  at  our  disposal.  Successful  precipitation  must  in- 
volve conditions  which  insure  the  nearest  approximation  to 
insolubility  of  the  precipitated  body,  and  the  precipitate  must 
also  be  in  a  form  favorable  for  filtration  and  washing.  For 
crystalline  precipitates,  the  latter  condition  is  fulfilled  when 
the  crystals  are  relatively  large.  This  is  often  attained  by 
allowing  the  fine  crystals,  which  first  separate,  to  digest  in 
contact  with  the  hot  liquid  from  which  they  have  fallen. 
During  this  digestion  the  smaller  crystals,  which  are  very 
slightly  more  soluble  than  the  larger  ones,  reclissolve,  and 
the  solution,  which  is  supersaturated  as  regards  the  larger 
crystals,  allows  the  latter  to  separate.  This  transfer  is  fur- 
ther promoted  by  the  influence  of  surface  tension,  which 
tends  to  reduce  the  surface  of  the  solid,  i.  e.,  to  increase 
the  size  of  the  individual  crystals. 

Certain  amorphous  bodies,  such  as  ferric  hydroxide,  alu- 
minum hydroxide,  and  silicic  acid  may  pass  into  a  colloidal 
state,  in  which  they  form  semi-solutions.  This  may  happen 
if  an  attempt  is  made  to  precipitate  them  from  solutions 
which  are  free  from  other  salts.  These  rarely  occur  in  anal- 
ysis, but  during  the  washing  of  these  precipitates  such  semi- 
solutions  may  sometimes  be  formed,  unless  some  salt  is 
added  to  the  wash  water.  In  cases  where  the  addition  of 
a  salt  to  the  wash  water  is  impracticable,  the  precipitate 
should  be  digested  for  some  time  on  the  steam  bath  with 
the  original  solution,  a  procedure  which  lessens  its  ten- 
dency to  pass  into  the  colloidal  state. 


FILTRATION.  !3 

In  all  precipitations  the  reagent  should  be  added  slowly, 
with  constant  stirring,  and  should  be  hot  when  circumstances 
permit.  The  slow  addition  is  less  likely  to  occasion  contam- 
ination of  the  precipitate  by  the  inclosure  of  other  substances 
which  may  be  in  the  solution,  or  of  the  reagent  itself. 

For  the  complete  removal  of  precipitates  from  containing 
vessels,  it  is  often  necessary  to  rub  the  sides  of  these  ves- 
sels to  loosen  the  adhering  particles.  This  can  best  be  done 
by  slipping  over  the  end  of  a  stirring  rod  a  piece  of  soft 
rubber  tubing,  which  has  been  well  washed  to  remove  loose 
fragments,  or  by  using  a  piece  of  sheet  rubber,  which  may 
be  folded  over  the  rod  and  cemented  together  by  moisten- 
ing the  surfaces  with  benzene.  The  sides  of  the  beaker  can 
then  be  rubbed  with  the  covered  rod. 

All  stirring  rods  should  have  the  ends  rounded  in  the 
flame  to  avoid  scratching  the  beakers. 


FILTRATION,     AND     THE     TESTING     OF     FILTRATES     AND 
WASHINGS. 

Distilled  water  should  be  employed  in  all  quantitative  work, 
and  nitration  should  be  made  only  through  "washed  filters," 
i.  e.,  those  which  have  been  treated  with  hydrochloric  and 
hydrofluoric  acids,  and  which,  on  incineration,  leave  a  small 
and  definitely  known  weight  of  ash.  Such  filters  are  read- 
ily obtainable  in  the  market. 

Funnels  should  be  selected  which  have  an  angle  as  near 
60°  as  possible,  and  with  a  narrow  stem  about  six  inches  in 
length.  The  filter  should  be  accurately  folded  to  fit  the 
funnel,  and  placed  so  that  the  top  of  the  filter  is  about  one- 
fourth  inch  below  the  top  of  the  funnel.  Under  no  cir- 
cumstances should  the  filter  ever  extend  above  the  edge  of 
the  funnel,  as  it  is  then  utterly  impossible  to  effect  com- 
plete washing. 

To  test  the  efficiency  of  the  filter,  fill  it  with  distilled 
water ;  this  water  should  soon  fill  the  neck  completely, 
forming  a  continuous  column  of  liquid  which,  by  its  hydro- 
static pressure,  produces  a  gentle  suction,  materially  pro- 
moting the  rapidity  of  filtration.  Unless  the  filter  allows 


I4  GENERAL    DIRECTIONS. 

a  free  passage  of  water  under  these  conditions,  its  use  is 
likely  to  prove  a  source  of  annoyance. 

The  use  of  a  vacuum  pump  to  promote  filtration  is  rarely 
altogether  advantageous  in  quantitative  analysis,  if  paper  fil- 
ters are  employed.  The  tendency  of  precipitates  to  pass 
through  the  pores  of  the  filter  is  increased,  and  this  source 
of  danger  more  than  compensates  for  the  possible  gain  in 
time.  Exception  may  be  made  in  the  case  of  such  precip- 
itates as  the  hydroxides  of  chromium,  aluminum,  or  iron,  and 
of  silicic  acid,  but  whenever  suction  is  applied,  the  point  of 
the  paper  filter  must  be  supported  by  a  perforated  platinum 
cone  or  a  small  "  hardened  filter  "  of  parchment.  The  rate 
of  filtration  is  often  greater  when  an  asbestos  felt  (Gooch 
filter)  is  used  (see  page  27  for  a  description),  and  the  pos- 
sibility of  a  substitution  of  this  for  the  paper  filter  should 
always  be  considered. 

When  the  filtrate  is  received  in  a  beaker,  the  stem  of  the 
funnel  should  touch  the  side  of  the  receiving  vessel  to  avoid 
loss  by  spattering.  Neglect  of  this  precaution  is  a  frequent 
source  of  error. 

The  vessels  which  contain  the  initial  filtrate  should  always 
be  replaced  by  clean  ones,  properly  labelled,  before  the  wash- 
ing of  a  precipitate  begins.  In  many  instances  a  finely  di- 
vided precipitate,  which  shows  no  tendency  to  pass  through 
the  filter  at  first,  while  the  solution  is  relatively  dense,  ap- 
pears at  once  in  the  washings.  Under  such  conditions  the 
advantages  accruing  from  the  removal  of  the  first  filtrate 
are  obvious,  both  as  regards  the  diminished  volume  requir- 
ing refiltration,  and  also  the  lesser  amount  of  washing  sub- 
sequently required. 

Much  time  may  often  be  saved  by  washing  precipitates 
by  decantation,  /'.  e.,  by  pouring  over  them,  while  still  in  the 
original  vessel,  considerable  volumes  of  wash-water  and  al- 
lowing them  to  settle.  The  supernatant,  clear  wash-water 
is  then  decanted  through  the  filter,  so  far  as  is  practicable 
without  disturbing  the  precipitate,  and  a  new  portion  of 
wash-water  is  added.  This  procedure  can  be  employed  to 
special  advantage  with  gelatinous  precipitates,  which  fill  up 
the  pores  of  the  filter  paper.  As  the  medium  from  which 


IGNITION  OF  PRECIPITATES.  15 

the  precipitate  is  to  settle  becomes  less  dense,  it  subsides 
less  readily,  and  it  becomes  necessary  to  transfer  it  to  the 
filter  and  complete  the  washing  there. 

A  precipitate  should  never  fill  the  filter  completely,  and 
the  wash-water  should  be  applied  at  the  top  of  the  filter, 
above  the  precipitate.  It  may  be  shown  mathematically 
that  the  washing  is  most  rapidly  accomplished  by  filling  the 
filter  well  to  the  top  with  wash-water  each  time,  and  allow- 
ing it  to  drain  completely  after  each  addition,  but  that  when 
a  precipitate  is  to  be  washed  with  the  least  possible  volume 
of  liquid  the  latter  should  be  applied  in  repeated  small  quan- 
tities. For  a  discussion  of  this  matter  and  the  phenomena 
of  adsorption,  the  student  is  referred  to  Ostwald's  Founda- 
tions of  Analytical  Chemistry,  page  15,  et  seq. 

Gelatinous  precipitates  should  not  be  allowed  to  dry  be- 
fore complete  removal  of  foreign  matter  is  effected.  They 
are  likely  to  shrink  and  crack,  and  subsequent  additions  of 
wash-water  pass  through  these  channels  only. 

Solutions  should  be  filtered  while  hot,  as  far  as  possible, 
since  the  motion  of  the  liquid  through  the  pores  of  a  filter 
is  retarded  by  internal  friction,  and  this,  for  water  at  100° 
C.,  is  less  than  one  sixth  of  the  resistance  at  o°  C. 

All  filtrates  and  wash-waters  without  exception  should  be 
properly  tested.  In  testing  the  latter  an  amount  not  less 
than  3  cc.  should  be  taken  for  the  final  test. 

It  is  impossible  to  trust  to  one's  judgment  with  regard  to 
the  washing  of  precipitates  ;  the  washings  from  each  pre- 
cipitate of  a  series  simultaneously  treated  must  be  tested, 
since  the  rate  of  washing  will  often  differ  materially  under 
apparently  similar  conditions.  No  exception  can  ever  be  made 
to  this  rule. 

The  habit  of  placing  a  clean  filter  paper  under  the  receiv- 
ing beaker  during  filtration,  is  one  to  be  commended.  On 
tnis  paper  a  record  of  the  number  of  washings  can  very  well 
be  made  as  the  portions  of  wash-water  are  added. 

IGNITION    OF    PRECIPITATES. 

The  larger  number  of  precipitates  may,  if  proper  precau- 
tions are  taken,  be  ignited  without  previous  drying.  If,  how- 


X6  GENERAL    DIRECTIONS, 

ever,  such  precipitates  can  be  dried  without  loss  of  time  to 
the  analyst  (as,  for  example,  over  night),  it  is  well  to  submit 
them  to  this  process.  It  should,  nevertheless,  be  remem- 
bered that  a  partially  dried  precipitate  requires  as  much,  or 
more  care  during  ignition  than  a  thoroughly  moist  one. 

The  precipitate,  with  the  filter  folded  over  it,  should  be 
placed  well  at  the  base  of  the  crucible,  which  should  then 
be  placed  so  far  above  the  flame  of  the  lamp  that  no  vio- 
lent escape  of  steam  is  possible.  When  the  filter  and  con- 
tents have  dried,  the  crucible  should  be  placed  on  its  side 
without  the  cover,  and  the  heat  should  be  gently  increased 
until  the  filter  chars,  but  should  never  be  increased  beyond 
this  point  until  all  volatile  matter  from  the  dry  distillation 
of  the  filter  paper  has  been  expelled  without  taking  fire. 
Much  annoyance  will  be  avoided  by  observing  this  point. 

During  this  preliminary  heating  the  flame  should  be  placed 
near  the  mouth  of  the  crucible,  but  in  all  subsequent  heat- 
ing the  flame  of  the  lamp  should  be  well  at  the  base  of  the 
crucible,  as  it  is  inclined  upon  its  side,  to  allow  a  ready 
access  of  oxygen  and  to  avoid  the  entrance  of  unburned 
(reducing)  gases.  When  the  filter  has  charred,  the  heat 
should  be  raised  to  redness  until  ignition  is  complete. 
The  heating  of  precipitates  over  the  blast  lamp  is  to  be 
avoided  unless  specially  directed. 

The  limited  number  of  instances  in  which  the  precipitate 
must  be  separated  from  the  filter  preliminary  to  ignition 
will  be  treated  of  as  they  occur. 

USE    AND    CARE    OF    BALANCES. 

The  analytical  balance  is  a  delicate  instrument,  which  will 
perform  excellent  service  under  careful  treatment,  but  such 
treatment  is  an  essential  condition  if  its  accuracy  is  to  be 
depended  upon.  The  following  rules  may  be  regarded  as 
embodying  the  important  points  involved  in  the  use  of  a  bal- 
ance, but  no  rules  can  do.  away  with  the  necessity  for  a 
sense  of  personal  responsibility  on  the  part  of  each  student, 
since  by  carelessness  he  can  render  inaccurate  not  only  his 
own  analyses,  but  those  of  other  students  using  the  balance  : 


USE   OF  BALANCES.  ^ 

1.  The  balance  pans  should  be  brushed  off  and  the  ad- 
justment of  the  balance  tested  before  use,  particularly  where 
several  persons  use  the  same  instrument. 

To  determine  whether  or  not  the  balance  is  in  adjust- 
ment, note  (i)  whether  it  is  level;  (2)  whether  the  pointer 
rests  at  zero  when  the  beam  is  lifted  from  its  knife-edges, 
and  also  when  lowered  so  that  the  pan  arrests  touch  the 
scale  pans ;  (3)  that  the  mechanism  for  raising  and  lowering 
the  beams  works  smoothly ;  (4)  that  the  pan  arrests  touch 
the  pans  when  the  beam  is  lowered  ;  and  (5)  that  the  needle 
swings  equal  distances  on  either  side  of  the  zero-point  when 
set  in  motion  without  any  load  on  the  pans.  If  the  latter 
condition  is  not  absolutely  fulfilled,  the  balance  should  be 
adjusted,  unless  the  variation  is  not  more  than  one  division 
on  the  scale  ;  it  is  often  better  to  make  a  proper  allowance 
for  this  zero  error  rather  than  to  disturb  the  balance  by  an 
attempt  at  correction. 

Unless  a  student  thoroughly  understands  the  construction 
of  a  balance  he  should  never  attempt  to  make  adjustments, 
but  should  apply  to  the  instructor  in  charge.  For  a  dis- 
cussion of  the  construction  and  essential  characteristics  of 
a  balance  the  student  is  referred  to  Fresenius  Quantitative 
Analysis. 

2.  The  beam  should  never  be  set  in  motion  by  lowering 
it  forcibly  upon  the  knife-edges,  nor  by  touching  the  pans, 
but  rather  by  means   of  the   rider   (unless  the   balance  be 
provided   with    some    of    the    newer   devices   for    the    pur- 
pose),  and   the    swing   should    be    arrested    only  when    the 
needle  passes  the  zero  on  the   scale,  otherwise  the  knife- 
edges    become    dull.      For    the    same    reason    the     beam 
should    never    be    left    upon    its    knife-edges,    nor    should 
weights  be  removed   from,  or  placed  on  the  pans  without 
supporting  the  beam,  except  in  the  case  of  the  small  frac- 
tional weights. 

3.  In  testing  the  weight  of  a  body,  the  weights  should  be 
applied  in  the  order  in  which  they  occur  in  the  weight-box 
(not  at  haphazard),  and  the  weight  should  be  recorded  first 
by  noting  the  weights   missing   from   the  weight-box,   and 
that    record    subsequently   checked    as    these   weights   are 


1 8  GENERAL    DIRECTIONS. 

taken    from    the   pan.      This   practice   will   often    avoid   or 
detect  errors. 

4.  The  balance-case  should  always  be  closed  during  the 
final  weighing,  when  the  rider  is  used,  to  protect  the  pans 
from  the  influence  of  air  currents. 

Before  the  final  determination  of  an  exact  weight  the 
beam  should  always  be  lifted  from  the  knife-edges  and 
again  lowered  into  place,  as  it  frequently  happens  that 
the  scale  pans  are  twisted  by  the  impact  of  the  weights,  the 
beam  being  virtually  lengthened  or  shortened.  Lifting  the 
beam  restores  the  proper  alignment. 

After  the  weighing  is  finished,  the  weights  should  always 
be  replaced  in  their  proper  places  in  the  weight-box  and 
the  rider  taken  from  the  beam. 

5.  No  chemical  substance  should  ever  be  placed  directly 
upon  the  balance-pan.     Every  substance  or  vessel  weighed 
should  be  dry  and  cold.     A  warm  object  occasions  the  for- 
mation of  air  currents,   which  vitiate  the  accuracy  of  the 
weight. 

6.  Above  all,  if  any  damage  be  done  to  a  balance,  if  any 
substance  be  spilled  upon  the  pans,   or  if  the  mechanism 
appear  to  be  deranged,   the   matter   should   receive   imme- 
diate attention,  and  should  be  reported  at  once  to  the  in- 
structor in   charge.      In  the  majority  of  instances   serious 
damage  can  be  averted  by  prompt  action,  when  delay  might 
ruin  the  balance. 


NOTEBOOKS. 

Notebooks  should  contain,  beside  the  record  of  observa- 
tions, descriptive  notes.  All  records  of  weights  should  be 
placed  upon  the  right-hand  page,  while  that  on  the  left  is 
reserved  for  the  notes,  calculations  of  factors,  or  the  amount 
of  reagents  required. 

The  neat  and  systematic  arrangement  of  the  records  of 
analyses  is  of  the  first  importance,  and  is  an  evidence  of 
careful  work  and  an  excellent  credential.  Of  two  notebooks 
in  which  the  results  may  be,  in  fact,  of  equal  value  as  legal 


ECONOMY  OF  TIME.  !9 

evidence,  that  one  which  is  neatly  arranged  will  carry  with 
it  greater  weight. 

All  records  should  be  dated,  and  all  observations  should 
be  recorded  at  once  in  the  notebook.  The  making  of  rec- 
ords upon  loose  paper  is  a  practice  to  be  deprecated,  as  is 
also  that  of  copying  original  entries  into  a  second  notebook. 
The  student  should  accustom  himself  to  orderly  entries  at 
the  time  of  observation. 

The  descriptive  notes  should  mention  any  special  diffi- 
culties encountered  in  the  analyses  and  the  remedies  applied,, 
and  also  incidents  in  the  course  of  the  analysis,  if  any,  which 
may  possibly  influence  the  results  injuriously.  All  analyses, 
should  be  made  in  duplicate,  and  in  general  a  close  agree- 
ment in  results  should  be  expected.  It  should,  however,  be 
remembered  that  a  close  concordance  of  results  in  "check 
analyses  "  is  not  conclusive  evidence  of  the  accuracy  of  those 
results,  although  the  probability  that  such  is  the  case  is,  of 
course,  considerably  enhanced.  The  satisfaction  in  obtain- 
ing "check  results"  in  such  analyses  must  never  be  allowed 
to  interfere  with  the  critical  examination  of  the  procedure 
employed,  nor  must  they  ever  be  regarded  as  in  any  meas- 
ure a  substitute  for  absolute  truth  and  accuracy. 

ECONOMY    OF    TIME. 

An  economical  use  of  laboratory  hours  is  best  secured  by 
acquiring  a  thorough  knowledge  of  the  character  of  the  work 
to  be  done  before  undertaking  it,  and  then  by  so  arranging 
the  work  that  no  time  shall  be  wasted  during  the  evapora- 
tion of  liquids  and  like  time-consuming  operations.  To  this 
end  the  student  should  read  thoughtfully  not  only  the  pro- 
cedure, but  the  explanatory  notes  as  well,  before  any  step  is 
taken  in  the  analysis. 

Several  analyses  should  be  in  progress  at  once  and  con- 
fusion carefully  guarded  against  by  a  free  use  of  labels. 

In  general,  economy  of  time  results  from  the  filtration  of 
several  solutions  at  once,  since  the  washing  of  five  or  more 
precipitates  may  frequently  be  accomplished  in  the  time  re- 
quisite for  any  one,  if  taken  alone. 


20  GENERAL    DIRECTIONS. 

ACCURACY   AND    INTEGRITY    DEMANDED. 

The  fundamental  conception  of  quantitative  analysis  im- 
plies a  necessity  for  all  possible  care  in  guarding  against 
loss  of  material,  or  the  introduction  of  foreign  matter.  All 
filters  and  solutions  should  be  covered  to  protect  them  from 
dust,  just  as  far  as  is  practicable,  and  every  particle  of  solu- 
tion or  precipitate  must  be  regarded  as  invaluable  for  the 
success  of  the  analysis. 

In  this  connection  it  must  also  be  emphasized  that  only 
the  operator  himself  can  know  the  whole  history  of  an  an- 
alysis, and  only  he  can  know  whether  his  work  is  worthy  of 
full  confidence.  No  work  should  be  continued  for  a  moment 
after  such  confidence  is  lost,  but  should  be  resolutely  dis- 
carded as  soon  as  a  cause  for  distrust  is  fully  established. 
The  student  should  determine  to  put  forth  his  best  efforts 
in  each  analysis  ;  it  is  well  not  to  be  too  ready  to  condone 
failures  and  to  "begin  again,"  as  much  time  is  lost  in  these 
fruitless  attempts.  Nothing  less  than  absolute  integrity  is  or 
can  be  demanded  of  a  quantitative  analyst,  and  arty  disregard 
of  this  principle,  however  slight,  is  as  fatal  to  success  as  lack 
of  chemical  knowledge  or  inaptitude  at  manipulation  can  pos- 
sibly be. 


PART   II. 


GRAVIMETRIC   ANALYSIS. 


DETERMINATION  OF  CHLORINE  IN  SODIUM  CHLORIDE. 


PREPARATION. 

THE  preparation  of  chemically  pure  sodium  chloride  from 
the  commercial  article  may  be  effected  as  follows  : 

Procedure. — Weigh  out,  upon  rough  balances,  about  50 
grams  of  a  sample  of  "  table  salt,"  cover  this  with  120  cc. 
of  distilled  water,  stir  until  the  water  is  saturated,  and  filter. 
To  the  filtrate  add  concentrated  hydrochloric  acid  (sp.  gr. 
1. 20)  until  the  chloride  begins  to  separate,  then  pass  into 
the  solution  gaseous  hydrochloric  acid.  This  acid  should  be 
generated  in  a  flask,  from  rock  salt  and  commercial  sulphu- 
ric acid.  The  gas  should  be  washed  by  passing  it  through 
concentrated,  aqueous  hydrochloric  acid,  and  the  delivery- 
tube  should  terminate  in  a  2-inch  funnel,  placed  mouth 
downward,  to  prevent  the  clogging  of  the  delivery-tube 
by  the  separated  salt.  When  the  separation  of  the  salt 
has  apparently  ceased,  remove  it  by  filtration  upon  a 
paper  disc  placed  upon  a  perforated  porcelain  plate  (a 
Witt  filter),  and  drain  by  suction.  Wash  the  chloride  with 
25  cc.  of  hydrochloric  acid  (sp.  gr.  1.12)  in  successive  small 
portions,  allowing  the  precipitate  to  drain  completely  after 
each  addition.  Wash  finally  with  a  small  quantity  (5  cc.)  of 
water,  and  test  this  wash-water  for  sulphates.  If  sulphates 
are  found,  the  washing  with  hydrochloric  acid  must  be  con- 
tinued. When  freed  from  sulphates,  transfer  the  precipitate 
to  a  porcelain  or  platinum  dish,  or  crucible,  and  heat  until 
decrepitation  ceases.  The  chloride  should  then  be  allowed 


22  GRAVIMETRIC   ANALYSIS. 

to  cool  in  a  desiccator,  and  be  placed  in  a  weighing-tube  (like 
a  small  test-tube),  which  should  be  kept  tightly  stoppered. 

Notes.  —  i.  The  commercial  grades  of  table  salt  contain, 
beside  sodium  chloride,  chlorides  or  sulphates  of  magnesium, 
calcium,  or  potassium,  the  two  first-named  causing  the  salt  to 
absorb  moisture.  When  hydrochloric  acid  is  added  to  a 
saturated  solution  of  the  salt,  sodium  chloride  is  thrown 
down,  leaving  the  impurities  in  solution.  The  principle  un- 
derlying this  separation  may  be  briefly  stated  by  saying 
that  the  solution  is  saturated  with  respect  to  chlorine  radi- 
cals (chlorine  ions,  more  properly),  and  when  more  of  these 
chlorine  ions  are  added,  in  the  form  of  the  readily  soluble 
hydrochloric  acid,  some  of  those  in  combination  with  sodium 
are  forced  out  of  solution  in  the  form  of  the  relatively  insol- 
uble sodium  chloride. 

2.  The  precipitation  of  the  sodium  chloride  might  be  ef- 
fected more  quickly  by  the  addition  of  liberal  quantities  of 
concentrated  aqueous  hydrochloric  acid,  but  its  purity  is  less 
certain  under  those  conditions.     The  slow  separation,  caused 
by  the  absorption  of  the  gas,  is  more  favorable  to  the  isola- 
tion of  a  pure  product,  and  the  process  is   also  somewhat 
more  economical. 

3.  Since  the  sodium   chloride   is   not  insoluble   in   either 
the  acid  or  the  water  used  for  washing,  it  is  essential  that 
these  should  be  used  in  as  small  quantities  as  is  practicable. 
Note  the  statement  on  page  15  concerning  the  most  efficient 
method  of  washing  a  precipitate  with  a  limited   quantity  of 
liquid. 

4.  The  heating  of  the  chloride  is  essential  to  expel  any 
excess  of  hydrochloric  acid  held  by  the  salt,  and  to  remove 
moisture  inclosed  between  crystal  surfaces.      The  escape  of 
this  moisture  is  the  cause  of  decrepitation.     Even  the  pure 
salt  is  slightly  hygroscopic  ;  hence  the  necessity  for  cooling 
in  the  dry  air  of  the  desiccator  and  for  preservation  in  stop- 
pered tubes. 

ANALYSIS. 

The  sodium  chloride,  prepared  as  above,  is  ready  for  analy- 
sis, and  if  the  preparation  has  been  carefully  made,  the  per- 
centage of  chlorine  found  on  analysis  should  agree  closely 
with  that  calculated  from  the  symbol. 


DETERMINATION   OF   CHLORINE. 


23 


Procedure,  (a)  Carefully  clean  the  weighing-tube  containing 
the  sodium  chloride,  handling  it  as  littje  as  possible  with  the 
moist  fingers,  and  weigh  it  accurately  to  o.oooi  gram,  re- 
cording the  weight  at  once  in  the  notebook.  Hold  the  tube 
over  the  top  of  a  No.  3  lipped  beaker,  and  cautiously  remove 
tfte  stopper,  noting  carefully  that  no  particles  fall  from  re, 
or  from  the  tube,  elsewhere  than  into  the  beaker.  Pour  out 
a  small  portion  of  the  chloride,  replace  the  stopper,  and  de- 
termine by  approximate  weighing  how  much  has  been  re- 
moved. Continue  this  procedure  until  0.25-0.30  gram  has 
been  taken  from  the  tube,  then  weigh  accurately  and  record 
the  weight  beneath  the  first  in  the  notebook.  The  differ- 
ence of  the  two  weights  represents  the  weight  of  the  chlo- 
ride taken  for  analysis.  Again  weigh  a  second  portion  of 
0.25-0.30  gram  into  a  second  beaker  of  the  same  size  as 
the  first.  The  beakers  should  be  plainly  marked  to  corre- 
spond with  the  entries  in  the  notebook.  Dissolve  each 
portion  of  the  chloride  in  150  cc.  of  distilled  water,  and  add 
about  ten  drops  of  nitric  acid  (sp.  gr.  1.20).  Calculate  the 
volume  of  a  silver  nitrate  solution  required  to  effect  com- 
plete precipitation  in  each  case,  and  add  slowly  about  5  cc. 
in  excess  of  that  amount,  with  constant  stirring.  Heat  the 
solutions  cautiously  to  boiling,  stirring  occasionally,  and  con- 
tinue the  heating  and  stirring  until  the  precipitates  settle 
promptly,  Jeaving  a  nearly  clear  supernatant  liquid.  This 
heating  should  not  take  place  in  direct  sunlight.  The 
beaker  should  be  covered  with  a  watch-glass,  and  both 
boiling  and  stirring  $o  regulated  as  to  preclude  any  possi- 
bility of  loss  of  material.  Add  to  the  clear  liquid  one  or 
two  drops  of  silver  nitrate  solution,  to  make  stire  that  an 
excess  of  the  reagent  is  present.  Prepare  two  washed  fil- 
ters (9  cm.  in  diameter),  bearing  in  mind  the  precautions 
mentioned  on  pages  13  and  14,  and  pass  the  liquid  through 
the  filter,  leaving  the  chloride  in  the  beaker  as  far  as  pos- 
sible. Wash  the  precipitates  two  or  three  times  with 
hot  water,  by  decantation,  transfer  them  to  the  filter  by 
means  of  a  stream  from  the  wash-bottle,  with  the  aid  of  a 
feather  or  a  rubber  tip  on  a  stirring  rod,  if  need  be.  Fi- 
nally remove  the  main  filtrate,  replace  by  a  clean  beaker, 


24  GRAVIMETRIC   ANALYSIS. 

and  wash  the  filters  and  precipitates  until  3  cc.  of  the  wash- 
ings show  no  cloudiness  with  a  drop  of  hydrochloric  acid. 
The  funnels  should  then  be  covered  with  a  filter  paper 
which  has  been  previously  moistened  and  stretched  over 
the  sides  of  the  funnel,  to  which  it  will  adhere  on  drying. 
It  should  be  properly  labelled  with  the  student's  name  and 
desk  number,  and  then  placed  in  a  drying  closet,  at  a  tem- 
perature of  about  100-110°  C.,  until  completely  dry.  Put 
the  filtrate  containing  the  silver  nitrate  aside  in  a  suitable 
receptacle  for  "silver  residues,"  from  which  the  silver  can 
be  recovered. 

The  perfectly  dry  filter  is  then  opened  over  a  circular 
piece  of  clean,  smooth,  glazed  paper  about  six  inches  in  di- 
ameter, placed  upon  a  larger  piece  about  twelve  inches  in 
diameter.  The  precipitate  is  removed  from  the  filter  as 
completely  as  possible,  by  rubbing  the  sides  gently  to- 
gether, or  by  scraping  them  cautiously  with  a  feather, 
which  has  been  cut  close  to  the  quill  and  is  slightly  stiff. 
In  either  case,  care  must  be  taken  not  to  rub  off  any  con- 
siderable quantity  of  the  paper,  nor  to  lose  silver  chloride 
in  the  form  of  dust.  Cover  the  precipitate  on  the  glazed 
paper  with  a  watch-glass,  to  prevent  loss  of  fine  particles 
and  to  protect  it  from  dust.  Fold  the  filter  paper  care- 
fully, as  it  was  when  it  came  from  the  funnel,  roll  it  into  a 
small  cone,  and  wind  loosely  around  the  top  a  piece  of  small 
platinum  wire.  Hold  the  filter  by  the  wire  over  the  proper 
porcelain  crucible,  ignite  it,  and  allow  the  ash  to  fall  into 
the  crucible.  Place  the  crucible  upon  a  clean  clay  triangle, 
on  its  side,  and  ignite,  with  the  flame  well  at  its  base  un- 
til all  the  carbon  of  the  filter  has  been  consumed.  Allow 
the  crucible  to  cool,  add  two  drops  of  nitric  acid  and  one 
drop  of  hydrochloric  acid,  and  heat  very  cautiously,  to  avoid 
spattering,  until  the  acids  have  been  expelled;  then  trans- - 
fer  the  main  portion  of  the  precipitate  from  the  glazed 
paper  to  the  cooled  crucible,  placing  the  latter,  for  the  pur- 
pose, on  the  larger  piece  of  glazed  paper  and  brushing  the 
precipitate  from  the  smaller  piece  into  it,  sweeping  up  all 
particles  belonging  to  the  determination. 

Moisten  the  precipitate  with  two  drops  of  nitric  acid  and 


DETERMINATION   OF    CHLORINE.  25 

one  drop  of  hydrochloric  acid,  and  again  heat  with  great 
caution  until  the  acids  are  expelled  and  the  precipitate  is 
white,  after  which  the  temperature  is  gradually  raised  until 
the  silver  chloride  begins  to  fuse.  The  crucible  is  then 
cooled  in  a  desiccator  and  weighed,  after  which  the  heat- 
ing (without  the  addition  of  acids)  is  repeated,  and  it  is 
again  weighed.  This  must  be  continued  until  the  weight 
is  constant  within  0.0003  gram  in  two  consecutive  weigh- 
ings. Deduct  the  weight  of  the  crucible,  and  calculate  the 
weight  of  chlorine  in  the  silver  chloride,  and  subsequently 
the  percentage  in  the  sample  of  sodium  chloride  taken  for 
analysis.  Consult  Part  V,  page  132. 

Notes.  —  i.  The  nitric  acid  is  added  before  precipitation 
to  lessen  the  tendency  of  the  silver  chloride  to  carry  down 
with  it  other  substances  which  may  be  present  in  the  solu- 
tion. A  large  excess  of  the  acid  would  exert  a  slight  solvent 
action  upon  the  chloride. 

2.  The  solution  should  not  be  boiled  after  the  addition  of 
the  nitric   acid,  before   the   presence  of  an   excess  of  silver 
nitrate  is  assured,  since  a  slight  interaction  between  the  nitric 
acid  and  the  sodium  chloride  is  possible,  by  which  a  loss  of 
chlorine,  either  as  such  or  as  hydrochloric  acid,  might  ensue. 
The  presence  of  an  excess  of  the  precipitant  can  usually  be 
recognized  at  the  time  of  its  addition,  by  the  increased  readi- 
ness with  which  the  precipitate  clots  together  and  settles. 

3.  The  precipitate  should  not  be  exposed  to  strong  sun- 
light, since  by   its  action  a  reduction  of  the  silver  chloride 
is  effected,  accompanied  by  a  loss  of  chlorine.     The  superfi- 
cial alteration  which  the  chloride  undergoes  in  diffused  day- 
light is  not  fatal  to  the  accuracy  of  the  determination,  since 
the  slight  loss  of  chlorine  may  be  counteracted  by  the  treat- 
ment of  the  precipitate  with  nitrohydrochloric  acid,  as  noted 
below. 

4.  The  precipitate  must  be  washed  with  hot  water  until  it 
is  absolutely  free  from  silver  and  sodium  nitrates.     It  may  be 
assumed  that  the  sodium  is  also  completely  removed  when  the 
wash-water  shows  no  evidence  of  silver.     It  must  be  borne 
in  mind  that  silver  chloride  is  somewhat  soluble  in  hydro- 
chloric acid,  and  only  a  single  drop  should  be  added.     The 


26  GRAVIMETRIC   ANALYSIS. 

washing  should   be   continued   until   no   cloudiness  whatever 
can  be  detected  in  3  cc.  of  the  washings,  v 

5.  The  separation  of  the  silver  chloride  from  the  filter  is 
essential,  since  the  burning  carbon  of  the  paper  would  reduce 
a  considerable  quantity  of  the  precipitate  to  metallic  silver, 
and  its  complete  reconversion  to  the  chloride  within  the  cru- 
cible, by  means  of  acids,  would  be  accompanied  by  some  un- 
certainty.     The  small   amount  of   precipitate  which  adheres 
to  the  filter  is  partially  reduced  to  metallic  silver  during  the 
ignition,  but  this  small  quantity  can  be  dissolved  in  the  nitric 
acid  which  is  added,  and  completely  reconverted  to  chloride 
by  the  hydrochloric  acid.     The  subsequent  addition  of  these 
two  acids  to  the  main  portion  of  the  precipitate  restores  the 
chlorine  to  the  chloride  reduced  by  the  sunlight. 

The  platinum  wire  is  wrapped  around  the  top  of  the  filter 
during  its  incineration,  to  avoid  contact  with  any  reduced  sil- 
ver which  might  come  from  the  reduction  of  the  precipitate. 
If  the  wire  was  placed  nearer  the  apex,  such  contact  could 
hardly  be  avoided. 

6.  Silver  chloride  should  not  be  heated  to  complete  fu- 
sion, since  a  slight  loss  by  volatilization  is  possible  at  higher 
temperatures.      The    temperature    of   fusion    is    not    always 
sufficient  to   destroy  filter  shreds ;    hence   these   should  not 
be  allowed  to  contaminate  the  precipitate. 

7.  The   ignited  precipitate   of   silver  chloride   should   be 
placed  in  the  jar  for  "silver  residues."      The  crucible  may 
he  cleaned  by  placing  in  it  some  granulated  zinc  and  dilute 
sulphuric  acid.     The  chloride  is  soon  loosened,  and  may  be 
detached. 

8.  Silver  chloride  is  practically  insoluble  in  water  and  di- 
lute nitric  acid,  slightly  soluble  in  strong  nitric  acid,  and  ap- 
preciably so  in  strong  hydrochloric  acid.     It  is  also  slightly 
soluble  in  hot  concentrated  solutions  of  silver  nitrate.     The 
chloride  is  readily  soluble  in  aqueous  ammonia  and  in  solu- 
tions of  potassium  cyanide  and  sodium  thiosulphate. 

9.  Stoppers   in    weighing-tubes    are    likely   to    change    in 
weight   from    the    varying    amounts    of    moisture    absorbed 
from  the  atmosphere.     It  is,  therefore,  necessary  to  confirm 
the  recorded  weight  of  a  tube  which   has  been  unused  for 
some  time,  before  weighing  out  a  new  portion  of  substance 
from  it. 


DETERMINATION   OF    CHLORINE. 


27 


DETERMINATION     OF     CHLORINE     IN     SODIUM     CHLORIDE,     WITH 
THE    USE    OF    A    GOOCH    FILTER. 

A  commercial  sample  of  table  salt  may  advantageously  be 
substituted  for  the  pure  sodium  chloride,  if  the  latter  has 
already  been  examined.  The  table  salt  should  be  heated 
until  decrepitation  ceases,  and  cooled  in  a  desiccator. 

Procedtire.  —  (b)  Weigh  out  two  portions  of  the  sub. 
stance,  each  weighing  about  0.25  gram,  and  precipitate 
the  silver  chloride  as  described  in  procedure  (a).  Mean- 
while prepare  a  Gooch  filter  as  follows :  Select  a  small 
glass  funnel,  i  to  i£  inches  in  diameter,  and  stretch  over 
its  mouth  a  piece  of  rubber-band  tubing  ("bill-tie  tubing") 
about  i  inch  wide  and  i|  inches  long.  This  should  be 
drawn  down  on  the  sides  of  the  funnel  until  it  holds  firmly, 
leaving  an  opening  at  the  center  of  the  mouth  of  the  funnel, 
into  which  a  perforated  porcelain  crucible  (Gooch  crucible) 
is  fitted.  The  rubber  should  be  drawn  up  around  the  sides 
of  the  crucible  until  it  is  air-tight.  Fit  the  glass  funnel 
into  the  stopper  of  a  filter  bottle,  and  connect  it  with  the 
vacuum  pump.  Suspend  some  finely  divided  asbestos,  which 
has  been  washed  with  acid,  in  20  to  30  cc.  of  water ;  allow 
this  to  settle,  pour  off  the  very  fine  particles,  and  then 
pour  the  rest  cautiously  into  the  crucible  until  an  even  felt 
of  asbestos,  not  over  ^  inch  in  thickness,  is  formed.  A 
gentle  suction  must  be  applied  while  preparing  this  felt. 
Wash  the  felt  thoroughly  with  distilled  water  until  all 
fine  or  loose  particles  are  removed,  then  place  the  cruci- 
ble in  a  small  beaker,  and  place  both  in  a  drying  closet  at 
130°  C.  for  30  to  40  minutes.  Cool  the  crucible  in  a  des- 
iccator, and  weigh.  Heat  again  for  20  to  30  minutes,  cool, 
and  again  weigh,  repeating  this  until  the  weight  is  constant 
within  0.0003  gram.  The  filter  is  then  ready  for  use. 

Replace  the  crucible  in  the  funnel,  and  apply  a  gentle 
suction,  after  which,  pour  in  the  solution  to  be  filtered,  with- 
out disturbing  the  asbestos  felt.  When  pouring  liquid  into 
a  Gooch  filter,  hold  the  stirring  rod  well  down  in  the  cru- 


28  .  GRAVIMETRIC  ANALYSIS. 

cible,  so  that  the  liquid  does  not  fall  with  any  force  upon 
the  asbestos. 

Transfer  the  whole  of  the  precipitate  to  the  filter,  and 
wash  thoroughly  with  hot  water  until  free  from  soluble  sil- 
ver salts,  then  dry,  at  130°  C,  to  a  constant  weight.  The 
percentage  of  chlorine  may  be  calculated  from  the  weight 
of  silver  chloride. 

Notes. —  i.  The  asbestos  should  be  of  the  finest  quality 
and  capable  of  division  into  minute  fibrous  particles.  A 
coarse  felt  is  not  satisfactory. 

The  use  of  the  Gooch  filter  commends  itself  strongly  when 
a  considerable  number  of  halogen  determinations  are  to  be 
made,  since  successive  portions  of  the  silver  halides  may  be 
filtered  on  the  same  filter,  without  the  removal  of  the  pre- 
ceding portions,  until  the  crucible  is  about  two  thirds  filled. 
The  use  of  a  perforated  disc  of  porcelain  or  platinum,  which 
may  be  placed  upon  the  top  of  the  asbestos  felt,  serves  to 
protect  it  in  some  measure,  but  it  is  obvious  that  care  should 
be  taken  to  avoid  loosening  the  felt  at  the  edges  as  the  liquid 
is  poured  upon  it.  If  the  felt  is  properly  prepared,  filtration 
and  washing  are  rapidly  accomplished  on  this  filter,  and  this 
factor,  combined  with  possibility  of  collecting  several  precip- 
itates on  the  same  filter,  are  strong  arguments  in  favor  of 
its  use  with  any  but  gelatinous  precipitates.  If  perforated 
platinum  crucibles  are  employed,  which  can  be  fitted  into  a 
platinum  cap  after  removal  from  the  funnel,  the  precipitates 
can  be  ignited  over  the  flame  of  a  lamp  as  in  an  ordinary 
platinum  crucible.  The  asbestos  is  apt  to  curl  away  from  the 
edges  during  such  heating,  and  if  the  same  filter  be  used  for 
a  second  time,  great  care  is  required  to  prevent  loss  of  asbestos. 

2.  A  funnel  tube,  made  from  stout  glass  tubing  about 
ij  inch  inside  diameter  and  3  inches  long,  to  which  is  at- 
tached a  tube  of  suitable  size  and  length  to  pass  through  a 
rubber  stopper,  may  be  substituted  for  the  glass  funnel 
above  prescribed.  It  is  then  only  necessary  to  cover  the 
upper  edges  of  this  tube  with  rubber.  The  crucible  may  be 
pressed  into  it,  and  makes  air  tight  connections. 


FERROUS   AMMONIUM  SULPHATE. 


29 


DETERMINATION    OF    IRON    AND    OF    SULPHUR    IN 
FERROUS    AMMONIUM    SULPHATE. 

FeSO4.  (NH4)2SO4.  6H2O. 

DETERMINATION    OF    IRON. 

Procedure.  —  Select  a  quantity  of  perfectly  clear  crystals 
of  the  salt  sufficient  to  fill  a  weighing  tube.  -Weigh  out, 
into  two  No.  4  lipped  beakers,  two  portions  of  about  I  gram 
each,  and  dissolve  these  in  50  cc.  of  water,  to  which  i  cc. 
of  hydrochloric  acid  (sp.  gr.  1.12)  has  been  added.  Heat  the 
solution  to  boiling,  and  while  at  the  boiling  point  add  nitric 
acid  (sp.  gr.  1.42),  drop  by  drop  (noting  the  volume  used), 
until  the  brOwn  coloration,  which  appears  after  the  addition 
of  a  part  of  the  nitric  acid,  gives  place  to  a  yellow  or  red. 
Avoid  a  large  excess  of  nitric  acid,  but  be  sure  that  the  ac- 
tion is  complete.  Pour  this  solution  cautiously  into  about 
200  cc.  of  water,  containing  a  slight  excess  of  ammonia. 
Calculate  for  this  purpose  the  amount  of  aqueous  ammonia 
required  to  neutralize  the  acids  added,  and  also  to  precipitate 
the  iron  as  ferric  hydroxide  from  the  weight  of  the  ferrous 
ammonium  sulphate  taken  for  analysis.  The  volume  thus 
calculated  will  be  a  slight  excess  over  that  actually  required, 
since  a  part  of  the  acids  are  consumed  in  the  oxidation  proc- 
ess, or  are  volatilized.  Heat  the  solution  to  boiling,  and 
allow  the  precipitated  ferric  hydroxide  to  settle.  Decant 
the  clear  liquid  through  a  washed/filter  (9  cm.),  keeping  as 
much  of  the  precipitate  in  the  Weaker  as  possible.  Wash 
once  by  decantation  with  100  cc.  of  hot  water,  and  then 
transfer  the  bulk  of  the  precipitate  to  the  filter.  Dissolve 
the  iron  from  the  filter  with  hot  hydrochloric  acid,  and  col- 
lect the  solution  in  the  beaker  in  which  precipitation  took 
place.  A.dd  3  cc.  of  nitric  acid  (sp.  gr.  1.42),  boil  for  a  few 
moments,  and  again  pour  into  an  excess  of  ammonia.  Wash 
the  precipitate  twice  by  decantation,  and  finally  throw  it  on 
the  filter,  and  wash  continuously  with  hot  water  until  3  cc. 
of  the  washings  show  no  evidences  of  the  presence  of  chlo- 
rides when  tested  with  silver  nitrate,  acidified  with  nitric 


3° 


GRAVIMETRIC  ANALYSIS. 


acid.      The  filtrate  and  washings  are  combined  with  those 
from  the  first  precipitation. 

The  moist  precipitate  is  placed  in  a  platinum  crucible 
which  has  been  previously  heated,  cooled  in  a  desiccator, 
and  weighed.  It  is  then  treated  according  to  the  directions 
for  "Ignition  of  Precipitates,"  page  16.  When  the  volatile 
matter  of  the  filter  has  been  expelled,  raise  the  temperature 
to  the  full  heat  of  the  burner  for  fifteen  minutes,  and  finally 
heat  over  the  blast  lamp,  with  the  crucible  covered,  for  three 
minutes.  Cool  and  weigh.  Repeat  the  strong  heating 
until  the  weight  is  constant  within  0.0003  gram.  Exercise 
great  care  when  heating  over  the  blast  lamp  that  a  small 
flame  is  used,  and  that  this  is  directed  against  the  bottom 
of  the  crucible  in  such  a  way  as  to  preclude  the  entrance 
of  unburned  or  reducing  gases  into  it,  by  reflection  from 
the  edges  of  the  cover.  From  the  weight  of  ferric  oxide 
(Fe2O3)  calculate  the  weight  of  iron  (Fe)  and  the  percent- 
age of  the  latter  in  the  sample. 

Notes.  —  i.  If  a  selection  of  pure  material  for  analysis  is 
to  be  made,  those  crystals  which  are  cloudy  are  to  be  avoided, 
on  account  of  loss  of  water  of  crystallization  ;  and  also  those 
which  are  red,  indicating  the  presence  of  ferric  iron.  If,  on 
the  other  hand,  the  value  of  an  average  sample  of  the  material 
is  desired,  it  is  preferable  to  grind  the  whole  together,  mix 
thoroughly,  and  take  from  the  mixture  a  sample  for  analysis. 

2.  The  hydrochloric  acid  is  added  to  prevent  the  precipi- 
tation of  basic  ferric  salts  during  solution,  as  a  result  of  a  par- 
tial oxidation  of  the  iron  in  the  absence  of  free  acid.     The 
nitric  acid  oxidizes  the  ferrous  iron,  after  attaining  a  mod- 
erate strength,  with  the  formation  of  an  intermediate  nitroso- 
compound  similar  in  character  to  that  formed  in  the  "  ring- 
test"  for  nitrates.      The  nitric  oxide  is  driven  out  by  heat, 
and  the  solution  then  shows  by  its  color  the  presence  of  ferric 
chloride.     A  drop  of  the  oxidized  solution  may  be  tested  on  a 
watch-glass  with  potassium  ferricyanide,  to  insure  the  absence 
of  ferrous  salts.     This  oxidation  of  the  iron  is  necessary,  since 
ferrous  salts  are  not  completely  precipitated  by  ammonia. 

3.  The  ferric  hydroxide   tends  to  carry  down   some  sul- 
phuric acid  in  the  form  of  basic  ferric  sulphate.     This  ten- 


DETERMINATION   OF  IRON.  31 

dency  is  lessened  if  the  solution  of  the  iron  is  added  to  an 
excess  of  ammonia,  since  under  these  conditions  immediate 
and  complete  precipitation  of  the  hydroxide  ensues  ;  whereas, 
by  the  gradual  neutralization  with  ammonia,  the  opportunity 
for  the  local  formation  of  a  neutral  solution  within  the  liquid, 
and  consequent  deposition  of  a  basic  sulphate,  is  favored. 
Even  with  this  precaution  the  entire  absence  of  sulphates 
from  the  first  iron  precipitate  is  not  assured.  It  is,  there- 
fore, redissolved  and  again  thrown  down  by  ammonia.  The 
organic  matter  of  the  filter  paper  may  occasion  a  partial  re- 
duction of  the  iron  during  solution,  with  consequent  possibil- 
ity of  incomplete  precipitation  with  ammonia.  The  nitric 
acid  is  added  to  reoxidize  this  iron. 

4.  By  the  ignition  of  ferric  oxide  with   ammonium   chlo- 
ride, volatile  ferric  chloride  is  formed,  with  consequent  loss 
of    iron.      The    precipitate    must,    therefore,    be    completely 
washed.    The  washings  are  acidified  with  nitric  acid,  before 
testing  with  silver  nitrate,  to  destroy  the  ammonia,  which  is 
a  solvent  of  silver  chloride. 

The  use  of  suction  to  promote  filtration  and  washing  is 
permissible,  though  not  prescribed.  The  precipitate  should 
not  be  allowed  to  dry  during  the  washing,  for  reasons  stated 
on  page  15. 

5.  To  avoid   errors    arising   from    the   solvent   action    of 
ammonia'cal  liquids  upon  glass,  the   iron   precipitate  should 
be  filtered  without  unnecessary  delay. 

6.  The  directions  for  the  ignition  of  precipitates  must  be 
closely   followed.      A  ready  access  of   atmospheric   oxygen 
is  of  special  importance,  to  insure  the  reoxidation  to  ferric 
oxide  of  any  iron  which  may  be  reduced  to  magnetic  oxide 
during  the  combustion  of  the  filter.     The  final  heating  over 
the  blast  lamp  is  essential  for  the  complete  expulsion  of  the 
last  traces  of  water  from  the  hydroxide. 

7.  Ignited  ferric  oxide  is  somewhat  hygroscopic,  on  which 
account  the  weighings  must  be  promptly  completed  after  re- 
moval from  the  desiccator.      In  all  weighings  after  the  first, 
it  is  well  to  place  the  weights  upon  the  balance  pan  before 
removing  the  crucible  from  the  desiccator.     It  is  then  only 
necessary  to  move  the  rider  to  obtain  the  weight. 

8.  Ferric  hydroxide  is  practically  insoluble  in  ammoniacal 
liquids,  in  the  presence  of  ammonium   salts,  but  the   corre- 


32  GRAVIMETRIC  ANALYSIS. 

spending  hydroxides  of  aluminum  and  chromium  are  partially 
redissolved  by  an  excess  of  ammonia.  Chromium  hydroxide 
is  much  the  most  soluble  of  the  three.  In  other  respects 
the  gravimetric  determination  of  these  two  metals  is. compar- 
able with  that  of  iron. 

For  a  further  statement  of  the  properties  of  these  bodies 
the  student  is  referred  to  Frestnius1  Quantitative  Analysis, 
under  "Forms." 

DETERMINATION    OF    SULPHUR. 

Procedure.  —  Add  to  the  combined  filtrates  from  the  ferric 
hydroxide,  hydrochloric  acid  in  moderate  excess,  and  evapo- 
rate to  dryness  on  the  water  bath.  Add  10  cc.  of  hydro- 
chloric acid  (sp.  gr.  1.12)  to  the  residue,  and  again  evaporate 
to  dryness  on  the  bath.  Dissolve  the  residue  in  water,  fil- 
ter if  not  clear,  transfer  to  a  No.  5  lipped  beaker,  dilute  to 
about  400  cc.,  and  cautiously  add  hydrochloric  acid  until  the 
solution  shows  a  distinctly  acid  reaction.  Heat  the  solution 
to  boiling,  and  add  very  slowly,  and  with  constant  stirring, 
20  cc.  in  excess  of  the  calculated  amount  of  hot  barium  chlo- 
ride solution  (which  should  contain  about  20  grams  BaCl2, 
2H2O  per  liter).  Continue  the  boiling  for  about  two  min- 
utes, allow  the  precipitate  to  settle,  and  decant  the  liquid  at 
the  end  of  a  half  hour.  Replace  the  beaker  containing  the 
original  filtrate  by  a  clean  beaker,  wash  the  precipitated 
sulphate  by  decantation  with  hot  water,  and  subsequently 
upon  the  filter,  until  it  is  freed  from  chlorides.  The  filter  is 
then  transferred  to  a  platinum  crucible  and  ignited,  as  de- 
scribed on  page  16,  until  the  weight  is  constant. 

To  test  the  purity  of  the  precipitates,  mix  each,  in  the  cru- 
cible, with  five  to  six  times  its  weight  of  sodium  carbonate. 
This  can  best  be  done  by  placing  the  crucible  on  a  piece  of 
glazed  paper  and  stirring  the  mixture  with  a  clean,  dry  stir- 
ring rod,  which  may  finally  be  wiped  off  with  a  small  frag- 
ment of  filter  paper,  the  latter  being  placed  in  the  crucible. 
Cover  the  crucible  and  heat  over  a  Bunsen  or  Tirrill  burner 
until  a  quiet  liquid  fusion  ensues.  As  the  fused  mass  cools, 
insert  in  it  a  piece  of  platinum  wire,  coiled  so  that  it  will 
hold  securely  in  the  solidified  mass.  When  solidification  is 


DETERMINATION   OF   SULPHUR.  33 

complete,  replace  the  lamp  under  the  crucible  and  heat  only 
long  enough  to  cause  the  outside  of  the  mass  to  fuse.  Now 
allow  the  crucible  to  cool  completely,  when  the  mass  may 
frequently  be  at  once  drawn  out  of  the  crucible  by  the  wire. 
If  it  still  adheres,  a  cubic  centimeter  or  so  of  water  may  be 
placed  in  the  cold  crucible  and  cautiously  brought  to  boiling, 
when  the  cake  will  become  loosened,  and  may  be  removed  on 
the  wire  and  suspended  in  about  250  cc.  of  distilled  water 
to  dissolve. 

Extract  the  residue  of  barium  carbonate  thoroughly  with 
water,  taking  care  to  clean  the  crucible  completely.  Filter 
off  the  carbonate,  and  wash  it  with  hot  water,  testing  the 
washings  for  sulphate,  and  preserving  any  precipitates 
which  appear  in  these  tests.  Acidify  the  filtrate  with  hy- 
drochloric acid  until  just  acid,  bring  to  boiling,  and  add  hot 
barium  chloride  solution  slowly,  as  before.  Add  also  any 
tests  from  the  washings  in  which  precipitates  have  appeared. 
Filter,  wash,  ignite,  and  weigh.  Compare  the  results  with 
those  first  obtained,  and  calculate  the  weight  of  sulphur  in 
the  barium  sulphate,  and  from  that  the  percentage  in  the 
ferrous  ammonium  sulphate. 

Notes.  —  i.  Barium  sulphate,  in  a  larger  measure  than 
most  substances,  tends  to  carry  down  other  bodies  which  are 
present  in  the  solution  from  which  it  separates,  even  when 
these  other  bodies  are  relatively  soluble.  This  is  notably 
true  in  the  case  of  nitrates  and  chlorates  of  the  alkalies,  and 
of  iron,  and,  since  in  this  analysis  ammonium  nitrate  has  re- 
sulted from  the  neutralization  of  the  excess  of  nitric  acid 
added  to  oxidize  the  iron,  it  is  essential  that  this  should 
be  destroyed  by.  repeated  evaporation  with  a  relatively  large 
quantity  of  hydrochloric  acid.  During  evaporation  a  mutual 
decomposition  of  the  two  acids  takes  place,  and  the  nitric 
acid  is  finally  decomposed  and  expelled  by  the  excess  of 
hydrochloric  acid. 

2.  The  precipitation  of  the  sulphur  as  barium  sulphate 
might  take  place  in  the  presence  of  the  iron,  but  under  these 
conditions  the  likelihood  of  contamination  of  the  sulphate  by 
iron  would  be  considerable,  and  a  purification  of  the  precip- 
itate would  be  unavoidable.  On  the  other  hand,  ferric  chlo- 


34  GRAVIMETRIC  ANALYSIS. 

ride  exerts  a  slight  solvent  action  upon  the  barium  sulphate. 
For  these  reasons  it  is  preferable  to  precipitate  the  iron  be- 
fore proceeding  to  the  determination  of  the  sulphur. 

3.  Barium  sulphate  is  slightly  soluble  in  hydrochloric  acid, 
even  dilute ;  hence  only  the  smallest  excess  should  be  added 
over  the  amount  required  to  acidify  the   solution.      Recent 
investigations  show  that  the  presence  of  an  excess  of  the  ba- 
rium  chloride   lessens   the   solubility  of  the   sulphate  in   the 
acid.     An  addition  of  20  cc.  of  solution  in  excess  is,  there- 
fore, prescribed,  but,  for  the  reasons  stated  in  Note   i,  this 
excess  of  chloride  should  not  be  too  large,  and  for  the  same 
reasons  the  reagent  should  be  added  very  slowly,  and  with 
constant  stirring.     It  has  been  shown  that  the  rapid  addition 
leads  to  a  slight  co-precipitation  of  the  chloride,  which  can- 
not be  washed  out  of  the  sulphate. 

4.  The  precipitation  of  the  barium  sulphate   is  probably 
complete   at   the   end  of  a  half-hour,  and   the   solution   may 
safely  be  filtered  at  the  expiration  of  that  time,  if  it  is  desired 
to  hasten  the  analysis. 

As  noted  on  page  12,  many  precipitates  of  the  general  char- 
acter of  this  sulphate  tend  to  grow  more  coarsely  granular  if 
digested  for  some  time  with  the  liquid  from  which  they  have 
separated.  It  is,  therefore,  well  to  allow  the  precipitate  to  stand 
in  a  warm  place  for  several  hours  before  filtration,  whenever 
practicable,  to  promote  ease  of  filtration.  The  filtrate  and 
washings  should,  however,  always  be  carefully  examined  for 
minute  quantities  of  the  sulphate  which  may  pass  through 
the  pores  of  the  filter.  This  is  best  accomplished  by  impart- 
ing to  the  liquid  a  gentle  rotary  motion,  when  the  sulphate, 
if  present,  will  collect  in  the  center  of  the  beaker.  All  fil- 
trates in  this,  and  other  determinations,  must  be  tested  for 
complete  precipitation,  by  adding  to  them  a  small  quantity  of 
the  reagent  and  allowing  them  to  stand. 

5.  A  reduction  of  barium  sulphate  to  the  sulphide  may  be 
caused  by  the  reducing  action  of  the  burning  carbon  of  the 
filter,  but  subsequent  ignition,  with   ready  access  of  air,  re- 
converts the  sulphide  to  sulphate,  unless  a  considerable  reduc- 
tion has  occurred.     In  the  latter  case  it  is  expedient  to  add 
one  or  two  drops  of  sulphuric  acid,  and  to   heat  cautiously 
until  the  excess  of  acid  is  expelled. 

6.  Most  impurities   which   are    inclosed   by  the    sulphate 


DETERMINATION'  OF  SULPHUR. 


35 


cannot  be  removed  by  washing  with  water;  treatment  with 
hydrochloric  acid,  even  if  it  accomplishes  the  removal  of  these 
impurities,  dissolves  some  of  the  sulphate,  which  must  be  re- 
covered.  It  is  advisable,  then,  in  any  case,  and  essential  when 
the  contamination  is  due  to  iron,  alumina,  or  silica,  to  purify 
by  fusion  with  sodium  carbonate,  as  described  in  the  proce- 
dure. By  this  process  the  impurities  are  either  rendered  in- 
soluble, and  are  removed  by  filtration  with  the  barium  carbon- 
ate, or,  if  they  pass  into  solution  with  the  sodium  sulphate,, 
they  are  present  in  such  small  amounts  relatively,  that  they 
fail  to  be  cnrried  down  by  a  second  precipitation  of  the  sul- 
phate. It  is  obvious  that  the  excess  of  alkaline  carbonate 
must  be  destroyed  by  hydrochloric  acid,  and  that  the  same 
care  must  be  taken  in  the  addition  of  the  barium  chloride 
the  second  time,  as  was  taken  at  first.  The  reaction  during 
fusion  is  the  following : 

BaSO4  +  Na2CO3  =  Na2SO,  +  BaCO3. 

7.  The  removal  of  the  fused  mass  from  the  crucible  is  fa- 
cilitated by  the  procedure  outlined,  because,  after  the  second 
short  heating,  the  crucible,  by  its  more  rapid  cooling,  springs 
away  from  the  mass  inside.  The  boiling  with  water  is  some- 
times necessary  to  dissolve  a  slight  ring  of  carbonate,  which 
solders  the  mass  to  the  crucible  at  its  upper  edge. 

For  a  further  statement  of  the  properties  of  barium  sul- 
phate, the  student  is  referred  to  Fresenius1  Quantitative  Analy- 
sis, under  "  Forms." 


36  GRAVIMETRIC  ANALYSIS. 

DETERMINATION    OF   PHOSPHORIC   ANHYDRIDE   IN 

APATITE. 

The  sample  of  apatite  selected  for  analysis  should  be  as 
nearly  pure  as  possible.  Specimens  of  the  mineral  which 
leave  but  a  slight  siliceous  residue  are  not  difficult  to  secure. 

Procedure.  —  Grind  the  mineral  in  an  agate  mortar  until 
no  grit  is  perceptible.  Transfer  the  substance  to  a  weigh- 
ing tube,  and  weigh  out  two  portions,  not  exceeding  0.20 
gram  each,  into  two  No.  2  lipped  beakers.  Pour  over  them 
20  cc.  of  nitric  acid  (sp.  gr.  1.2),  and  warm  gently  until  sol- 
vent action  has  apparently  ceased.  Unless  the  absence  of 
soluble  silicates  is  assured,  evaporate  the  solution  cautiously 
to  dryness,  heat  the  residue,  for  about  two  hours,  to  130°  C., 
and  treat  it  again  with  nitric  acid,  as  described  above  ;  sep- 
arate the  residue  of  silica  by  filtration  on  a  small  filter  (7  cm.) 
and  wash  with  warm  water,  using  as  little  as  possible  (see  p. 
.  1 5).  Receive  the  filtrate  in  a  No.  3  or  No.  4  lipped  beaker. 
Test  the  washings  with  ammonia  for  calcium  phosphate,  but 
add  all  such  tests,  in  which  a  precipitate  appears,  to  the 
original  filtrate.  The  filtrate  and  washings  should  not  ex- 
ceed 100  cc.  in  volume.  Add  aqueous  ammonia  until  the 
precipitate  of  calcium  phosphate  first  produced  just  fails  to 
redissolve,  and  then  add  a  few  drops  of  nitric  acid  until 
this  is  again  brought  into  solution.  Warm  the  solution  until 
it  cannot  be  comfortably  held  in  the  hand,  and,  after  re- 
moval from  «the  lamp,  add  75  cc.  of  amnftonium  molybdate 
solution  (68  grams  MoO3  per  liter),  which  has  been  gently 
warmed,  but  which  must  be  perfectly  clear.  Allow  the  mix- 
ture to  stand  at  a  temperature  of  about  50°  to  60°  C.  for 
twelve  hours.  Filter  off  the  yellow  precipitate  on  an  9  cm. 
filter,  and  wash  by  decantation  with  a  solution  of  ammonium 
nitrate  made  acid  with  nitric  acid.*  Allow  the  precipitate 
to  remain  in  the  beaker  as  far  as  possible.  Test  the  wash- 
ings for  calcium  with  ammonia  and  ammonium  oxalate. 


*  This  solution  is  prepared  as  follows:  Mix  100  cc.  of  ammonia  solution 
(sp.  gr.  0.96)  with  325  cc.  of  nitric  acid  (sp.  gr.  1.2),  and  dilute  with  100  cc.  of 
water. 


DETERMINATION  OF  PHOSPHORIC  ANHYDRIDE.        37 

Add  10  cc.  of  molybdate  solution  to  the  filtrate,  and  leave 
it  for  a  few  hours.  It  should  then  be  carefully  examined 
for  a  yellozu  precipitate ;  a  white  precipitate  may  be  neg- 
lected. The  filtrate  should  not  be  thrown  away,  but  should 
be  placed  in  a  suitable  receptacle  for  "  molybdenum  resi- 
dues," from  which  the  molybdic  acid  may  be  recovered. 

Dissolve  the  precipitate  upon  the  filter,  by  pouring  through 
it  dilute  aqueous  ammonia  (one  volume  of  ammonia  (sp.  gr. 
0.96)  and  three  volumes  water,  which  should  be  carefully 
measured],  and  receive  the  .solution  in  the  beaker  containing 
the  bulk  of  the  precipitate.  The  total  volume  of  filtrate 
and  washings  must  not  exceed  100  cc. 

Calculate  the  volume  of  magnesium  ammonium  chloride 
solution  ("magnesia  mixture")  required  to  throw  out  the 
phosphoric  acid,  assuming  40  per  cent.  P2O5  in  the  apatite. 
Measure  out  not  more  than  2  cc.  in  excess,  and  add  this 
quantity  to  the  cold  ammoniacal  solution,  by  dropping  it 
from  a  glass  tube,  stirring  the  solution  constantly.  The 
rate  of  addition  must  not  be  greater  than  10  cc.  in  a  min- 
ute. Continue  the  stirring  for  a  few  moments,  and  set  the 
solution  aside,  at  the  temperature  of  the  laboratory,  over 
night.  (Or  it  may  be  stirred  constantly  for  a  half-hour, 
when  the  precipitation  should^  be  complete.)  The  magne- 
sium ammonium  phosphate  is  'then  removed  by  filtration  and 
washed  with  a  mixture  of  one  part  ammonia  (sp.  gr.  0.96), 
one  part  alcohol,  and  three  parts  water,  until  3  cc.  of  the 
washings  show  no  evidence  of  chlorides^  The  washings 
must  be  acidified  with  nitric  acid  before  the  silver  nitrate 
is  added.  Test  the  filtrate  carefully  for  complete  precipi- 
tation by  adding  more  magnesia  mixture  and  allowing  it  to 
stand. 

Cover  the  funnel  with  a  paper,  dry  the  filter  completely 
in  the  drying  closet,  and  then  ignite,  using  great  care  to 
insure  the  presence  of  plenty  of  oxygen  during  the  com- 
bustion of  the  filter  paper,  thus  guarding  against  a  possible 
reduction  of  the  phosphate,  with  disastrous  consequences 
both  to  crucible  and  analysis.  Do  not  raise  the  tempera- 
ture above  moderate  redness  until  the  precipitate  is  white. 
(Keep  this  precaution  well  in  mind.)  Ignite  finally  at  the 


38  GRAVIMETRIC  ANALYSIS. 

highest  temperature  of  the  Tirrill  burner,  and  repeat  the 
heating  until  the  weight  is  constant.  From  the  weight  of 
magnesium  pyrophosphate  (Mg2P2O7)  obtained,  calculate  the 
weight  of  phosphoric  anhydride  (P2O5),  and  the  percentage 
of  the  latter  in  the  sample  of  apatite. 

Notes. —  i.  Apatite  may  contain,  beside  calcium  phos- 
phate, either  calcium  fluoride  or  chloride.  It  is  evident  that 
the  direct  precipitation  of  the  phosphoric  acid  in  combination 
with  magnesium  is  impracticable  in  the  presence  of  any  metal 
which  forms  compounds  with  phosphoric  acid  which  are  insol- 
uble in  ammoniacal  liquids;  such,  for  example,  as'iron,  alumi- 
num, chromium,  and  the  alkaline  earths.  The  previous  isola- 
tion of  the  phosphoric  acid  in  combination  with  molybdenum, 
which  can  be  effected  in  nitric  acid  solution,  is  then  neces- 
sary. The  phospho-molybdate  is  soluble  in  ammonia,  and 
from  this  solution  the  phosphoric  acid  may  be  separated  as 
magnesium  ammonium  phosphate. 

2.  As  a  result  of  the  slight  solubility  of  magnesium  am- 
monium phosphate,  as  noted  below,  the  unavoidable  errors  of 
analysis'  are  greater  in  this  determination  than  in  those  which 
have  preceded  it,  and  some  divergence  may  be  expected  in 
duplicate  analyses.     It  is  obvious  that  the  larger  the  amount 
of  substance  taken  for  analysis,  the  less  will  be  the  relative 
loss  or  gain  due  to  experimental  errors ;  but  in  this  instance  a 
check  is  placed  upon  the  amount  of  material  which  may  be 
taken,  both  by  the  bulk  of   the  resulting  precipitate  of  am- 
monium phospho-molybdate,  and  by  the  excessive  amount  of 
ammonium  molybdate  required  to  effect  complete  separation 
of  the  phosphoric  acid,  since  a  liberal  excess  above  the  theo- 
retical quantity  is    demanded.     Molybdic  acid  is  one  of  the 
more  expensive  reagents. 

3.  Soluble  silicic  acid  might,  if  present,  partially  separate 
with    the    phospho-molybdate,    although    not  in   combination 
with  molybdenum.     Its  previous   removal  by  dehydration  is 
therefore  advisable. 

4.  Nitric  acid  is  chosen  as  a  solvent  because  the  phospho- 
molybdate  is  slightly  soluble  in  hydrochloric  acid.     An  excess 
of  nitric  acid  also  exerts  a  slight  solvent  action,  while  ammo- 
nium nitrate  lessens  the  solubility;  hence  the  neutralization 
of  the  former  by  ammonia. 


\/ 


DETERMINATION  OF  PHOSPHORIC  ANHYDRIDE.         39 

5.  The  composition  of  the  ''yellow  precipitate"  undoubt- 
edly varies  slightly  with  varying  conditions  at  the  time  of  its 
formation,   and   on   this   account   the   precipitate   is  not  com- 
monly separated  and  weighed   as  such.      Its  structure    may 
probably  be  represented  by  the  symbol,  (NH4)3PO4.  12  MoO3. 
H2O,  when  precipitated   under  the  conditions   prescribed  in 
the  procedure.     Whatever  other  variations  may  occuiitin  its 
composition,  the  ratio  of   12  MoO3  :  i  P  seems  to  holc^  and 
this  fact  is  utilized  in  volumetric  processes  for  the  determina- 
tion of  phosphorus,  in  which  the  molybdenum  is  reduced  to  a 
lower  oxide  and  re-oxidiaed  by  a  standard  solution  of  potas- 
sium permanganate. 

6.  The  precipitation  of  the  phospho-molybdate  takes  place 
moie  promptly  in  warm  than  in  cold  solutions,  but  the  temper- 
ature should  not  exceed -60°  C.  during  precipitation  ;  a  higher 
temperature  tends  to  separate  molybdic  acid  from  the  solu- 
tion.     This   acid   is  nearly  white,  and  its  deposition   in   the 
filtrate  on  long  standing  should  not  be  mistaken  for  a  second 
precipitation  of  the  yellow  precipitate. 

The   addition   of  75   cc.  of  ammonium  molybdate  solution 
insures  the  presence  of  a  liberal  excess  of  the  reagent. 

7.  When  washing  the  siliceous  residue,  the.  filtrate  may  be 
tested  for  calcium  by  adding  ammonia  alone,  since  that  re- 
agent neutralizes  the  acid  which  holds  the  calcium  phosphate 
in  solution  and  causes  precipitation  ;  but  after  the  removal  of 
the  phosphoric  acid  in  combination  with  the  molybdenum,  the 
addition  of  an  oxalate  is  required  to   show  the   presence  of 
calcium. 

8.  Magnesium  ammonium  phosphate  is  not  a  wholly  insol- 
uble body,  even  under  the   most  favorable  analytical  condi- 
tions.    It  is  least  soluble  in  a  liquid  containing  one  fourth  of 
its  volume  of  aqueous  ammonia  (sp.  gr.  0.96),  and  this  pro- 
portion should  be   carefully  preserved,   as  prescribed  in  the 
procedure.     On  account  of  this  slight  solubility,  the  volume 
of  solutions   should   be   kept   as   small   as  possible,  arid  the 
amount   of   wash-water    limited   to    that   absolutely  required. 
The   addition  of  alcohol  to  the  wash-water  lessens  the  sol- 
ubility of  the  magnesium  compound  in  it. 

9.  A  large  excess  of  the  magnesium  solution  tends  both  to 
throw  out  magnesium  hydroxide  (shown  by  a  flocculent  pre- 
cipitate), and  to  cause  the  phosphate  to  carry  down  molybdic 


4o  GRAVIMETRIC  ANALYSIS. 

acid.  The  latter,  if  its  presence  be  suspected,  may  be  re- 
moved from  the  phosphate  by  dissolving  the  precipitate  in 
hydrochloric  acid,  and  passing  sulphuretted  hydrogen  through 
the  warm  solution  from  three  to  four  hours. 

The  tendency  of  the  magnesium  precipitate  to  carry  down 
molybdic  acid  is  greater  if  the  solution  is  too  concentrated. 
The  volume  should  not  be  less  than  90  cc.  nor  more  than 
125  cc.  at  the  time  of  precipitation  with  the  magnesia  mixture. 

10.  The  magnesium  ammonium  phosphate  should  be  per- 
fectly crystalline,  and  will  be  so  if  the  directions  are  followed. 
The  slow  addition  of  the  reagent  is  essential,  and  the  stirring 
not  less  so.     Stirring  promotes  the  separation  of  the  precip- 
itate and  the  formation  of  larger  crystals,  and  may  therefore 
be   substituted  for  digestion  in  the  cold.      The  stirring  rod 
must  not  be   allowed   to  scratch  the  glass,   as    the    crystals 
adhere  to  such  scratches  and  are  removed  with  difficulty. 

The  remarks  on  page  12,  regarding  the  formation  of  large 
crystals  by  digestion  with  the  solution,  have  peculiar  force  in 
connection  with  the  magnesium  ammonium  phosphate,  which 
is  a  relatively  soluble  body. 

11.  During  ignition,  the  magnesium  ammonium  phosphate 
loses  ammonia  and  water,  and  is  converted  into  magnesium 
pyrophosphate. 

2  NH4MgPO4  =  Mg2P2O7  +  2  NH3  -+-  H2O. 

The  precautions  mentioned  on  page  16  must  be  observed 

/    with  great  care  during  the  ignition  of  this  precipitate.     The 

danger  here  lies  in  a  possible   reduction   of  the   phosphate. 

The  phosphorus  then   attacks  and  injures  the  crucible,  and 

the  determination  is  valueless. 

If  extreme  care  is  employed,  it  is  possible  to  safely  ignite 
this  precipitate  without  previous  drying,  but  the  student  is 
not  advised  to  attempt  this  until  some  general  experience  has 
been  gained. 

For  a  further  statement  of  the  properties  of  ammonium 
phospho-molybdate,  magnesium  ammonium  phosphate,  and 
magnesium  pyrophosphate,  the  student  is  referred  to  Fre- 
senius*  Quantitative  Analysis,  under  "  Forms." 


DOLOMITE.  4I 

DETERMINATION     OF     CALCIUM,    MAGNESIUM, 
AND    CARBON    DIOXIDE   IN    DOLOMITE. 

The  sample  of  dolomite  chosen  for  practice  should  leave 
little  or  no  residue  insoluble  in  hydrochloric  acid. 

DETERMINATION    OF    CALCIUM. 

Procedure.  —  Grind  the  mineral  to  a  fine  powder.  Weigh 
out,  into  1 50  cc.  casseroles,  two  portions  (a  and  b]  of  about 
one  gram  each.  Pour  over  them  30  cc.  of  hydrochloric  acid 
(sp.  gr.  1.12),  first  having  covered  the  casseroles;  add  i  cc. 
of  nitric  acid  (sp.  gr.  1.20)  and  boil  five  minutes.  Filter  off 
the  residue  on  a  small  filter,  wash  five  times  with  hot  water, 
and  ignite  the  filter  in  a  platinum  crucible.  Cover  the  filter 
ash  and  residue  with  a  small  quantity  (perhaps  0.5  gram)  of 
sodium  carbonate,  heat  to  fusion,  and  disintegrate  with  a 
little  water,  by  cautiously  boiling  it  in  the  crucible.  Add 
this  solution,  with  any  suspended  matter,  to  the  main  filtrate 
from  the  insoluble  residue;  evaporate  the  solution  to ' dry- 
ness,  and  heat  the  residue  in  the  hot  closet  for  two  hours 
at  130°  C.,  to  dehydrate  any  soluble  silicic  acid.  Moisten 
the  residue  with  hydrochloric  acid,  warm  gently,  dilute  to 
100  cc.,  and  bring  to  boiling.  Add  5  cc.  of  ammonium 
chloride  solution,  aqueous  ammonia  in  slight  excess,  and, 
if  manganese  is  present,  add  5  cc.  of  bromine  water  and  boil 
for  five  minutes,  adding  more  ammonia,  if  need  be.  If  man- 
ganese is  absent,  omit  the  bromine.  In  either  case,  finally 
neutralize  the  excess  of  ammonia  until  a  faint  odor  only 
can  be  detected,  and  filter  off  the  silicic  anhydride,  together 
with  the  iron,  aluminum,  and  manganese  hydroxides.  The 
filtration  must  take  place  promptly  after  the  addition  of  the 
ammonia.  Without  washing  the  precipitate,  re-dissolve  it 
on  the  filter  in  hydrochloric  acid,  receiving  the  solution  in  a 
clean  beaker.  Wash  the  filter  five  times  with  water,  and 
throw  down  the  hydroxides  from  the  solution  as  before. 
Wash  this  precipitate  with  hot  water  until  free  from  chlo- 
rides. The  precipitate  may  then  be  discarded. 

Treat  the  combined  filtrates  from  portion  (a)  as  follows  : 
To  the  filtrate,  concentrated  to  250  cc.  and  heated  to  boiling, 


42  GRAVIMETRIC  ANALYSIS. 

add  ammonium  oxalate  solution  in  moderate  excess,  stirring 
well,  and  adding  the  reagent  slowly.  Boil  for  two  minutes, 
allow  the  precipitated  calcium  oxalate  to  settle  for  a  half-hour, 
and  decant  through  a  filter.  Test  the  filtrate  for  complete 
precipitation  by  adding  a  few  drops  of  the  precipitant.  If  no 
precipitate  forms,  make  the  solution  slightly  acid  with  hydro- 
chloric acid,  and  proceed  with  the  magnesium  determination. 

Re-dissolve  the  calcium  oxalate  in  the  beaker,  and  from  the 
filter,  with  hydrochloric  acid,  washing  the  filter  five  times, 
and  finally  pouring  through  it  aqueous  ammonia.  Dilute  the 
solution  to  250  cc.,  bring  to  boiling,  add  I  cc.  ammonium 
oxalate  solution  and  ammonia  in  slight  excess  ;  boil  for  two 
minutes,  and  set  aside  for  a  half-hour.  Filter  off  the  cal- 
cium oxalate  upon  the  filter  first  used,  and  wash  free  from 
chlorides.  The  filtrate  should  be  made  barely  acid  and  com- 
bined with  the  first  filtrate. 

The  precipitate  of  calcium  oxalate  may  be  ignited  without 
drying.  It  should  be  ignited  at  the  highest  heat  of  the  Bun- 
sen  or  Tirrill  burner,  after  destroying  the  filter,  and  finally 
for  three  minutes  at  the  blast  lamp.  Repeat  until  the 
weight  is  constant.  As  the  calcium  oxide  absorbs  mois- 
ture from  the  air,  it  must  be  weighed  as  rapidly  as  possible. 
(Compare  Note  7,  on  p.  31.)  From  the  weight  of  the  oxide 
calculate  the  weight  of  calcium,  and  the  percentage  of  the 
"latter  in  the  dolomite. 

Treat  the  combined  filtrates  from  the  iron  precipitates  (ft) 
as  follows  :  Add  hydrochloric  acid  to  decided  acid  reaction. 
Calculate  the  volume  of  ammonium  oxalate  solution  required 
to  combine  with  both  the  calcium  and  magnesium,  assuming 
the  presence  of  20  per  cent,  of  each.  Add  this  volume  to 
the  solution  and,  if  a  precipitate  forms,  dissolve  it  in  hydro- 
chloric acid.  Dilute  the  whole  to  1000  cc.  and  bring  it  to 
boiling.  Remove  it  from  the  hot  plate  and  slowly  add 
ammonia  in  moderate  excess,  with  constant  stirring.  Allow 
the  solution  to  stand  for  a  half-hour,  filter  off  the  calcium 
oxalate,  and  wash  the  precipitate  with  hot  water  until  freed 
from  chlorides.  Ignite  this  precipitate  as  described  in  pro- 
cedure (a). 

Barely  acidify  the  filtrate  from  the  calcium  oxalate  with 


DETERMINATION  OF  MAGNESIUM.  43 

hydrochloric  acid,  and  proceed  with  the   magnesium   deter- 
mination. 

DETERMINATION    OF    MAGNESIUM. 

Procedure.  —  Evaporate  the  acidified  filtrates  from  the  cal- 
cium precipitates  until  salts  begin  to  crystallize,  .but  do  not 
evaporate  to  dryness.  Dilute  until  these  salts  are  brought 
into  solution.  Add  a  volume  of  aqueous  ammonia  (sp.  gr. 
0.96)  equal  to  one  third  the  volume  of  the  solution.  (To  do 
this,  measure  into  a  beaker  of  equal  size  a  volume  of  water 
equal  to  the  volume  of  the  solution.  Note  the  number  of 
cubic  centimeters  and  take  one  third  as  many  of  ammonia.) 
Calculate  the  volume  of  hydrogen  sodium  phosphate  solution 
required  to  precipitate  the  magnesium,  assuming  20  per 
cent,  to  be  present.  Add  the  sodium  phosphate  to  the  cold 
solution  drop  by  drop  from  a  glass  tube,  at  a  rate  not  greater 
than  10  cc.  per  minute;  stir  thoroughly,  set  aside  for  some 
hours,  and  treat  the  precipitate  of  magnesium  ammonium 
phosphate  as  prescribed  for  the  determination  of  phosphoric 
anhydride,  on  page  37. 

Notes.  —  i.  The  mineral  Dolomite  is  a  native  isomorphous 
mixture  of  calcium  and  magnesium  carbonates,  in  which  the 
relative  proportions  may  vary  widely.  Ferrous  iron  and  man- 
ganese are  not  infrequently  present,  and  most  specimens  leave 
a  larger  or  smaller  residue  on  treatment  with  hydrochloric 
acid.  Since  this  residue  may  contain  calcium  or  magnesium, 
it  must  be  rendered  soluble  by  fluxing,  and  brought  into  solu- 
tion. This  may  be  quickly  accomplished,  if  the  directions 
are  closely  followed  with  regard  to  the  use  of  a  small  quan- 
tity of  the  sodium  carbonate  and  disintegration  of  the  small 
fused  mass  within  the  crucible. 

If  it  is  definitely  known,  from  a  qualitative  examination, 
that  the  residue  insoluble  in  hydrochloric  acid  contains  no 
calcium  or  magnesium,  its  treatment  may  be  omitted,  and  the 
solution  may  be  at  once  evaporated  for  the  dehydration  of 
the  silica,  without  filtration.  The  insoluble  residue  may  be 
filtered  off  with  the  hydroxide  precipitate. 

2.  The  addition  of  nitric  acid  is  necessary  to  oxidize  fer- 
rous iron  to  ferric,  and  to  insure  its  complete  precipitation  by 


44 


GRAVIMETRIC  ANALYSIS. 

ammonia.  Manganese  is  not  oxidized  by  nitric  acid,  but  is 
oxidized  by  bromine  in  ammoniacal  solution  to  a  hydrated 
dioxide,  in  which  form  it  is  precipitated  and  removed  with 
the  iron  and  alumina.  The  possible  presence  of  alumina 
makes  it  necessary  to  finally  neutralize  all  but  a  slight  excess 
of  ammonia  before  filtration.  (Compare  Note  8,  p.  31.) 

3.  The   iron   precipitate   should  be  filtered  off  promptly, 
since  the  alkaline  solution  absorbs  carbon  dioxide  from  the 
air,  with  consequent  precipitation  of  the  calcium  as  carbon- 
ate.    This  is  possible  even  under  the  most  favorable  condi- 
tions, and  for  this  reason  the  iron  precipitate  is  re-dissolved 
and  thrown  out  again,  to  free  it  from  calcium. 

4.  The  separation  of  calcium  and  magnesium  as  oxalates 
requires   special   precautions.     It   is  necessary  either   to  re- 
dissolve  the  first  precipitate  of  calcium  oxalate,  and  make  a 
second  precipitation,  as  in  procedure  (a),  or  to  separate  in 
dilute  solution  and  in   the   presence  of  sufficient  ammonium 
oxalate  to  combine  with  both  metals,  as  in  procedure  (b).     In 
(a),    the   first   calcium    precipitate  contains   magnesium,    but 
when  dissolved,  the  proportion  of  magnesium  to  the  calcium 
in  solution  is  small,  and  the  second  calcium  precipitate  may 
be  considered  to  be  pure.     In  (£),  it  is  sought  to  accomplish 
the  same  end  by  liberal  dilution  in  the  presence  of  an  excess 
of  ammonium  oxalate,  lessening  in  this  way  the  tendency  of 
the  magnesium   to  separate  with  the   calcium.      Method  (a) 
probably  yields   results  which    are   more    certainly   accurate 
than  (£),  for  large  amounts  of  the   two  metals,  while  (b)  is 
accurate  for  small  quantities. 

5.  The  small  quantity  of  ammonium   oxalate   solution  is 
added  before  the  second  precipitation  of  the  calcium  oxalate, 
to  insure  the  presence  of  a  slight  excess  of  the  reagent,  which 
promotes  the  separation  of  the  calcium  compound. 

6.  On  ignition  the  calcium  oxalate  loses  carbon  dioxide 
and  carbon  monoxide,  leaving  calcium  oxide. 

CaC2O4  =  CaO  +  CO2  +  CO. 

For  small  weights  of  the  oxalate  (0.5  gram  or  less),  this 
reaction  may  be  brought  about  at  the  highest  temperature  of 
a  Tirrill  burner,  but  it  is  well  to  ignite  larger  quantities  than 
this  over  the  blast  lamp  until  the  weight  is  constant. 

7.  The  calcium  oxide  tends  to  absorb  moisture  to  form  the 
hydroxide  ;  hence  the  necessity  for  rapid  weighing. 


DETERMINATION'  £>F  CARBON  DIOXIDE.  45 

8.  The  filtrate  from  the  calcium  oxalate  should  be  made 
slightly  acid  immediately  after  filtration,  in  order  to  avoid  the 
solvent  action  of  the  alkaline  liquid  upon  the  glass. 

9.  The  precipitation  of  the  magnesium  should  be  made  in 
as  small  volume  as  possible,  and  the  ratio  of  ammonia  to  the 
total  volume   of  solution  should   be   carefully  provided   for, 
on  account  of  the  slight  solability  of  the  magnesium  ammo- 
nium phosphate.      This   matter  has  been  fully  discussed  in 
connection   with   the    phosphoric    anhydride    determination. 
(Compare  Note  8,  p.  39.) 

10.  If  the  magnesium  ammonium  phosphate  precipitate  is 
not  w.holly  crystalline,  as  it  should  be,  the  difficulty  may  some- 
times be   remedied  by  filtering  the  precipitate  and,  without 
washing  it,  re-dissolving  in  a  small  quantity  of  hydrochloric 
acid,  from  which  it  may  be  again  thrown  down  by  ammonia, 
after  adding  a  few  drops  of  sodium  phosphate  solution.     If 
the  flocculent  character  was  occasioned  by  the  presence  of 
magnesium  hydroxide,  the  second  precipitation,  in  a  smaller 
volume,  containing  fewer  salts,  will  often  result  more  favor- 
ably.    The  removal  of  iron  or  alumina  from  a  contaminated 
precipitate  is  a  matter  involving  a  long  procedure,  and  a  re- 
determination   of  the   magnesium   from   a  new  sample,   with 
additional  precautions,  is  usually  to  be  preferred. 

For  a  further  statement  of  the  properties  of  calcium  oxalate 
and  calcium  oxide,  the  student  is  referred  to  Fresenius'  Quan- 
titative Analysis,  under  "  Forms." 

DETERMINATION    OF    CARBON     DIOXIDE. 

The  apparatus  required  for  the  determination  of  the  car- 
bon dioxide  should  be  previously  arranged  as  shown  in  the 
cut  on  page  46.  The  flask  A  is  arranged  as  an  ordinary  wash 
bottle,  in  which  should  be  placed  50  cc.  of  hydrochloric  acid 
(sp.  gr.  1.12)  and  100  cc.  of  water.  This  flask  is  connected 
by  rubber  tubing  with  the  tube  (b)  leading  nearly  to  the 
bottom  of  the  evolution  flask  (B)  and  having  its  lower  end 
bent  upward  to  avoid  loss  of  gas  The  evolution  flask  may 
conveniently  be  a  wide-mouthed  Soxhlet  extraction  flask  of 
about  150  cc.  capacity,  the  mouth  of  which  is  fitted  with  a 
double-bored  rubber  stopper.  £7  is  a  ball-condenser  (or  may 
be  a  small  glass  [Liebig]  condenser),  from  the  top  of  which 


46  GRAVIMETRIC  ANALYSIS. 

a  delivery  tube  leads  to  the  U-tube  (D),  containing  some 
glass  beads  and  3  cc.  of  a  saturated  solution  of  silver  sul- 
phate with  3  cc.  of  concentrated  sulphuric  acid.  The  second 
tube  (E)  is  filled  with  calcium  chloride,  and  should  have  a 
small,  loose  plug  of  cotton  at  the  top  of  each  arm.  Both 
tubes  should  be  closed  by  cork  stoppers,  the  tops  of  which 
are  sunk  slightly  below  the  top  of  the  U-tube,  and  then 
neatly  sealed  with  sealing  wax. 

The  Geissler  bulb  (F)  should  be  so  filled  with  potassium 
hydroxide  solution  (sp.  gr.  1.27)  that  each  small  bulb  is 
about  two  thirds  full.  There  should  be  attached  to  the 
bulb,  in  such  a  manner  that  it  can  be  weighed  with  it,  a 
filled  calcium  chloride  tube  (3  in.).  A  platinum  wire  should 
be  attached  to  the  bulb  to  permit  it  to  hang  upon  the  support. 

An  additional  U-tube,  to  be  used  as  a  safety  tube  (H), 
should  be  filled  with  soda-lime  ready  for  use. 


When  the  apparatus  is  ready,  weigh  out  into  the  flask  (B) 
about  0.8  gram  of  the  dolomite  and  cover  it  with  15  cc.  of 
water.  Carefully  wipe  the  Geissler  bulb  (F),  with  its  cal- 
cium chloride  prolong  tube,  and  weigh  accurately.  After- 
ward stopper  the  openings  by  means  of  small  pieces  of 
rubber  tubing  closed  by  pieces  of  glass  rod. 


DETERMINATION'  OF  CARBON  DIOXIDE.  47 

Place  the  flask  (B)  on  the  apparatus,  making  sure  that 
the  stopper  fits  tightly,  and  connect  the  Geissler  bulb  with 
the  U-tubes  ;  disconnect  the  rubber  tube  from  A,  close  the 
pinch-cock  (a),  apply  suction  at  the  end  of  the  Geissler 
bulb,  and  note  whether  the  apparatus  is  absolutely  air-tight. 
This  precaution  must  not  be  neglected,  and  care  must  after- 
ward be  taken  to  admit  air  slowly,  by  cautiously  opening  (a), 
to  equalize  the  pressure.  The  rubber  tube  must,  of  course, 
contain  no  hydrochloric  acid  at  this  time. 

When  it  is  certain  that  the  apparatus  is  ready  for  use, 
connect  the  rubber  tube  with  the  flask  (A),  and,  opening  the 
pinch-cock  (a),  blow  over  about  10  cc.  of  the  diluted  acid. 
When  the  evolution  of  carbon  dioxide  slackens  add  a  fresh 
portion  of  acid.  The  gas  should  not  enter  the  Geissler 
bulb  more  rapidly  than  two  bubbles  per  second.  When  the 
action  of  the  acid  ceases  in  the  cold,  run  water  through  the 
condenser  (C)  and  apply  a  small  flame  to  the  flask,  cau- 
tiously bringing  the  liquid  to  the  boiling  point,  and  con- 
tinue to  boil  it  slowly  for  about  three  minutes.  Replace 
the  flask  (A)  by  the  safety  tube  (//).  Apply  suction  at  the 
end  of  the  Geissler  bulb,  and  regulate  it  in  such  a  way  that 
when  the  pinch-cock  (a)  is  opened  the  air  will  pass  slowly 
through  the  apparatus,  the  rate  not  to  exceed  that  named 
above.  Do  not  remove  the  lamp  from  under  the  flask  (B) 
until  the  air  current  is  adjusted.  Continue  to  draw  air 
through  the  apparatus  from  twenty  to  thirty  minutes ; 
then  disconnect  the  Geissler  bulb,  stopper  it,  place  it  in 
the  balance  room,  and,  after  an  interval  of  thirty  minutes, 
wipe  it  carefully,  and  weigh  it  without  the  stoppers.  The 
increase  in  weight  is  due  to  absorption  of  carbon  dioxide. 
Repeat  the  determination  with  a  fresh  portion  of  dolomite. 

Notes.  —  i.  By  very  cautious  procedure  it  is  possible  to 
avoid  the  use  of  a  condenser  above  the  evolution  flask,  or, 
a  tube  with  one  or  two  bulbs  blown  in  it  may  be  substi- 
tuted. More  moisture  is  then  likely  to  be  carried  over  into 
the  U-tubes,  which  soon  require  refilling. 

2.  The  air  current  may  conceivably  carry  minute  quan- 
tities of  hydrochloric  acid  through  the  apparatus.  This  acid 


48  GRAVIMETRIC  ANALYSIS. 

would  be  retained  by  the  silver  sulphate  solution  in  the  first 
U-tube.  The  beads  serve  to  divide  the  bubbles  of  gas  as 
they  pass  through  the  liquid.  The  sulphuric  acid  lessens  the 
evaporation  of  the  silver  sulphate  solution. 

The  calcium  chloride  in  the  second  tube  assures  the  re- 
moval of  the  water  from  the  moisture-laden  air  from  the 
evolution  flask.  As  calcium  chloride  frequently  contains 
basic  salts  which  would  retain  carbon  dioxide,  it  is  neces- 
sary to  pass  a  current  of  that  gas  through  the  U-tube  for  a 
short  time,  and  follow  this  by  a  current  of  dry  air  for  thirty 
minutes  before  using  the  tube. 

3.  The  potassium  hydroxide  unites  with  the  carbon  di- 
oxide to  form  potassium  carbonate.  The  hydroxide  must  be 
present  as  a  concentrated  solution  to  insure  complete  absorp- 
tion, and  to  avoid  a  loss  of  water  with  the  air  which  passes 
through  the  bulbs.  The  small  quantity  of  moisture  which  is 
taken  up  by  the  air  is  retained  by  the  calcium  chloride  of 
the  prolong  tube.  The  best  form  of  Geissler  bulb  is  that 
to  which  the  calcium  chloride  prolong  is  attached  by  a 
ground-glass  joint. 

5.  Carbon  dioxide  is  dissolved  by  cold  water,  but  the  gas 
is  expelled  by  boiling.     This  residual  volume  of  gas,  together 
with  that  which  is  distributed  through  the  apparatus,  is  swept 
out  through  the  absorption  bulb  by  the  current  of  air.     This 
is  purified   by  drawing  it  through  the   tube  (H)  containing 
soda-lime,  which  removes  any  carbon  dioxide  which  may  be 
in  the  atmosphere. 

6.  Instead  of  the  potassium  hydroxide  in  a  Geissler  bulb, 
soda-lime,  which  should  be  placed  in  two  U-tubes  with  suit- 
able  device  for  weighing,  may  be   employed  to   absorb  the 
carbon  dioxide.     Soda-lime  is  a  mixture  of  sodium  and  cal- 
cium hydroxides,  and  unites  with  the  gas  to  form  carbonates. 
Some  care  is  necessary  to  guard  against  loss  of  moisture  from 
the  soda-lime,  as  considerable  heat  is  generated  during  the 
absorption. 

7.  With  Geissler  bulbs  of  the  usual  size  two  determina- 
tions  may  be   safely  made  without  refilling  the    bulbs.      A 
precipitate   of  acid  potassium  carbonate   in  the  first  of  the 
bulbs  is  an  evidence  that  the  absorbing  power  of  the  hydrox. 
ide  is  nearly  exhausted. 

8.  Great  care  is  necessary  when  the  suction  is  applied  to 


DETERMINATION  OF  CARBON  DIOXIDE.  49 

avoid  a  violent  passage  of  gas  through  the  apparatus  and  to 
prevent  regurgitation. 

It  is  advisable  to  interpose  a  tube  (G)  containing  calcium 
chloride  and  soda-lime  between  the  suction  tube  and  the 
potash  bulb  to  prevent  possible  contamination. 

9.  The  large  surface  presented  to  the  air  by  the  Geissler 
bulbs  admits  of  the  accumulation  of  dust  and  moisture  dur- 
ing the  determination.     They  must,  therefore,  be  cautiously 
cleaned  before  each  weighing,  by  wiping  with  a  clean,  lintless 
cloth,   or  a  piece   of  wash-leather.      They  should  stand  for 
thirty  minutes  near  the  balance  to  assure  uniformity  of  tem- 
perature, any  change  of  which  at  the  second  weighing  of  so 
large  a  weight  might  cause  an  appreciable  error. 

The  stoppers  should  be  removed  from  the  bulbs  on  weigh- 
ing, that  the  air  inclosed  may  be  at  atmospheric  pressure. 

10.  A  large  number  of  procedures  have  been  proposed  for 
the  determination  of  carbon  dioxide  by  loss  of  weight,  but 
most  of  these  require  special  forms  of  apparatus.     Descrip- 
tions of  such  processes  can  be  found  in  any  large  work  on 
quantitative  analysis. 


50  GRAVIMETRIC  ANALYSIS. 

•  \~P  QJ\^~* 

DETERMINATION    OF    LEAD,    COPPER,     AND 
ZINC    IN    BRASS. 

DETERMINATION    OF    LEAD. 

Procedure.  —  Select  clean,  bright  chips  or  borings,  or,  if 
the  brass  is  in  the  form  of  wire,  polish  a  piece  of  suitable 
size  by  rubbing  with  emery,  cleaning  it  carefully  afterward. 
Weigh  out  two  portions  of  about  5  grams  each,  and  dissolve 
them  in  covered  casseroles  in  30  cc.  of  nitric  acid  (sp.  gr. 
1.2).  When  the  solution  is  complete,  cool,  wash  off  the 
cover  glass,  and  add  slowly  10  cc.  of  sulphuric  acid  (sp.  gr. 
1.84).  Evaporate  under  a  hood  until  heavy  white  fumes  of 
sulphuric  anhydride  are  evolved,  keeping  the  casserole  well 
covered  meanwhile;  cool,  add  125  cc.  of  water,  and  boil 
until  the  sulphates  of  copper  and  zinc  have  dissolved,  and  set 
the  solution  aside  until  perfectly  cold.  Filter  off  the  lead 
sulphate,  wash  it  by  decantation  with  dilute  sulphuric  acid 
(one  volume  of  concentrated  acid  to  twenty  volumes  of 
water)  until  the  washings  are  free  from  copper,  as  shown 
by  the  ammonia  test.  Set  the  filtrate  aside  for  the  deter- 
mination of  copper  and  zinc,  and  transfer  the  lead  sulphate 
to  the  filter,  finally  washing  the  filter  free  from  sulphuric 
acid,  with  alcohol  diluted  with  an  equal  volume  of  water. 
Be  sure  that  all  the  sulphate  is  removed  from  the  casserole. 
Discard  the  alcoholic  washings,  if  they  are  entirely  clear. 
Dry  the  filter,  and,  if  practicable,  separate  the  lead  sul- 
phate from  it,  as  described  on  page  24.  Prepare  a  No.  7 
porcelain  crucible  for  use,  by  heating  and  weighing,  burn 
the  filter  on  a  platinum  wire,  as  described  on  page  24,  and 
when  cold,  add  to  the  ash  two  drops  of  nitric  acid  and  one 
drop  of  sulphuric  acid.  Ignite  with  great  care  until  the 
acids  are  expelled.  Transfer  the  precipitate  which  was 
separated  from  the  filter,  to  the  crucible,  and  ignite  at  a 
moderate  red  heat.  Repeat  the  heating  until  the  weight  is 
constant. 

From  the  weight  of  the  lead  sulphate,  calculate  the  weight 
of  lead,  and  the  percentage  of  the  latter  in  the  brass. 


DETERMINATION  OF  LEAD.  51 

Notes. —  i.  It  is  obvious  that  the  brass  taken  for  analysis 
should  be  untarnished,  which  can  be  easily  assured  when  wire 
is  used,  by  scouring  with  emery.  If  chips  or  borings  are  used,, 
they  should  be  well  mixed,  and  the  sample  for  analysis  taken 
from  different  parts  of  the  mixture. 

2.  The  small  percentage  of  lead  usually  found  in  brasses 
makes  it  necessary  to  weigh  out  a  considerable  quantity  in 
order  to  secure  accuracy.     The  amount  taken,  5  grams,  is  too 
large  to  use  directly  for  the  determination  of  copper  and  zinc, 
on  account  of  the  bulk  of  the  precipitates  which  would  then 
have  to  be  handled.     An  aliquot  part  of  this  filtrate  is,  there- 
fore, used  for  these  analyses,  as  noted  under  the  determina- 
tion of  copper. 

3.  Lead  sulphate  is  slightly  soluble  in  nitric  acid;  hence 
the  latter  is  removed  by  heating  with  sulphuric  acid  until  the 
more  volatile  acid  is  expelled.     This  point  is  indicated  by  the 
appearance  of  the  heavy,  white  fumes. 

4.  Lead  sulphate  is  least  soluble  in  dilute  sulphuric  acid 
(i  :  20).     The  sulphuric  acid  solution  is,  therefore,  diluted  to 
secure  this  ratio,  and  a  wash-water  of  the  same  strength  is 
used. 

The  sulphuric  acid  of  the  wash-water,  if  allowed  to  remain 
on  the  filter,  would  char  it  during  the  drying,  making  subse- 
quent handling  difficult  or  impossible.  It  is  accordingly 
removed  by  washing  with  dilute  alcohol ;  but  the  alcohol  is  not 
added  to  the  main  filtrate,  as  its  presence  is  not  advantageous 
during  the  subsequent  operations. 

5.  The  lead  sulphate  must  be  separated  from  the  filter  for 
the  same  reasons  which  apply  in  the  case  of  silver  chloride, 
(Compare  p.  26.)     The  addition  of  nitric  acid  to  the  ash  dis- 
solves any  reduced  lead,  and  the  sulphuric  acid  converts  it  to- 
sulphate.     A  slight  loss  of  lead  is  possible  if  the  reduction  of 
any  considerable  quantity  of  the  precipitate  takes  place,  and 
it  is  evident  that  the  ignition  of  this  precipitate  in  platinum 
is  impracticable. 

6.  It  is  possible  to  determine  the  percentage  of  lead  in 
the  brass  by  solution  in  nitric  acid,  partial  neutralization  of 
the  excess  of  acid,  and  deposition  of  the  lead  by  an  electric 
current  as  peroxide  (PbO2)  on  the  anode ;  but  the  relatively 
large   percentage  of  copper  makes  the  simultaneous   deter- 
mination of  that  element  difficult,  if  a  sufficiently  large  quan- 


52  GRAVIMETRIC  ANALYSIS. 

tity  of  the  brass  be  taken  for  analysis  to  also  secure  an  accu- 
rate determination  of  the  lead.  For  the  determination  of 
small  quantities  of  lead  alone,  this  method  might  be  employed. 
For  a  further  statement  of  the  properties  of  lead  sulphate, 
the  student  is  referred  to  Fresenius'  Quantitative  Analysis, 
under  "  Forms." 

DETERMINATION    OF     COPPER. 

Procedure. —  Transfer  the  filtrate  from  the  lead  sulphate 
to  a  500  cc.  graduated  flask,  washing  out  the  beaker  care- 
fully. Fill  the  flask  with  distilled  water  until  the  lowest 
point  of  the  meniscus  is  exactly  level  with  the  mark  on  the 
neck  of  the  flask.  Carefully  remove,  with  a  strip  of  filter 
paper,  any  drops  of  water  which  are  on  the  inside  of  the 
neck  of  the  flask  above  the  graduation  ;  make  the  solution 
thoroughly  uniform  by  pouring  it  out  into  a  dry  beaker,  and 
back  into  the  flask  several  times,  and  finally  stopper  the 
flask.  Measure  off  one  tenth  of  this  solution  as  follows : 
Pour  into  a  50  cc.  graduated  flask  about  10  cc.  of  the  so- 
lution, shake  the  liquid  thoroughly  over  the  inner  surface 
of  the  small  flask  and  pour  it  out.  Repeat  the  same  op- 
eration. Fill  the  50  cc.  flask  until  the  lowest  point  of 
the  meniscus  is  exactly  level  with  the  mark  on  its  neck, 
remove  any  drops  of  solution  from  the  upper  part  of  the 
neck  with  filter  paper,  and  pour  the  solution  into  a  plain 
beaker  of  about  80  cc.  capacity,  of  rather  tall  form.  Wash 
out  the  flask  with  small  quantities  of  water  until  it  is  clean, 
adding  these  to  the  main  solution.  (Consult  also  note  on  p. 
72.)  When  the  second  portion  of  50  cc.  is  measured  out,  re- 
member that  the  flask  must  be  twice  rinsed  out  with  solution 
as  prescribed  above,  before  the  final  measurement  is  made. 

Add  to  the  copper  solution  i  gram  of  ammonium  nitrate. 
The  solution  is  then  ready  for  the  electrolytic  deposition  of 
the  copper. 

Meanwhile,  four  platinum  electrodes,  two  anodes  and  two 
cathodes,  should  be  cleaned  by  scouring  with  sapolio  and 
treatment  with  acid,  and  the  two  cathodes  ignited  gently  and 
cooled  in  a  desiccator.  Weigh  them  carefully,  and  place  one 
anode  and  one  cathode  in  each  solution.  The  connections 


DETERMINATION  OF  COPPER. 


55 


should  then  be  made  with  the  binding  posts  (or  other  device 
for  connection  with  the  electric  circuit),  in  such  a  way  that 
the  copper  will  be  deposited  upon  the  larger  electrode.  If  a 
gravity  battery  is  used,  this  electrode  should  be  connected 
with  the  zinc  pole ;  if  a  dynamo-current  is  used,  a  previous 
qualitative  experiment  is  necessary  to  determine  the  nega- 
tive pole.  Subject  the  solution  to  the  action  of  the  cur- 
rent from  eight  to  fifteen  hours,  as  indicated  by  the  rate  of 
deposition.  It  is  well  to  leave  the  solution  over  night,  if 
practicable. 

When  the  solution  is  colorless,  and  the  deposition  of  copper 
seems  to  be  complete,  break  the  circuit  as  follows :  Lift  the 
cathode  in  the  binding  post  until  it  is  nearly,  but  not  wholly, 
out  of  the  liquid,  and  wash  the  exposed  portion  with  distilled 
water.  Then  break  the  circuit  by  removing  the  electrode 
entirely,  and  complete  the  washing  promptly.  Pour  over 
the  electrode  and  copper  enough  alcohol  to  remove  adhering 
water,  and  dry  for  a  few  moments  at  105°  C.  .Cool  in  a  des- 
iccator and  weigh.  Test  the  solution  for  copper  by  the  addi- 
tion of  ammonia  in  excess,  to  about  5  cc.  of  the  solution.  If 
copper  is  found,  clean  the  electrodes  and  place  the  solution 
in  the  circuit  a  second  time.  Whether  copper  is  found  or 
not,  return  the  test  portion  to  the  main  solution. 

The  increase  in  weight  of  the  electrode  represents  directly 
the  weight  of  copper  from  one  tenth  of  the  solution,  from 
which  the  percentage  in  the  brass  may  be  calculated. 

Notes. —  i.  The  removal  of  one  tenth  of  the  solution  for 
the  determination  of  copper  and  zinc  is  necessitated  by  the 
difficulties  which  would  be  encountered  if  the  precipitation  of 
such  large  quantities  of  these  metals  as  would  come  from 
5  grams  of  the  brass,  was  attempted.  The  solution  is, 
therefore,  diluted  to  a  definite  volume  (500  cc.),  and  exactly 
one  tenth  (50  cc.)  is  measured  off.  To  attain  this  end 
it  is  evident  that  the  solution  must  be  thoroughly  uniform, 
which  is  brought  about  by  careful  mixing,  and  that  it  must 
not  be  diluted  by  water  adhering  to  the  small  flask.  This  is 
insured  by  rinsing  out  this  flask  several  times  with  the  solu- 
tion, instead  of  drying  it,  which  would  consume  much  time. 
The  two  flasks  must  be  graduated  at  the  same  temperature, 


GRAVIMETRIC  ANALYSIS. 

and  if,  as  is  usual,  the  small  flask  is  graduated  to  contain  50 
cc.,  it  must  be  rinsed  out  with  water  before  all  the  copper  and 
zinc  belonging  to  one  tenth  of  the  solution  are  obtained. 
This,  and  other  kindred  considerations,  are  discussed  under 
Volumetric  Analysis,  page  72.  A  pipette  might  properly  be 
substituted  for  the  small  flask,  if  desired. 

2.  The  presence  of  sufficient  ammonium  nitrate  to  react 
with  the  sulphuric  acid  in  the  solution,  lessens  the  tendency  of 
the  copper  to  deposit  on  the  cathode  in  a  spongy  condition. 
The  amount  of  the  nitrate  added  should  not  greatly  exceed 
i  gram. 

3.  The  electrodes  should  be  freed  from  all  greasy  matter 
before  using,  and  after  scouring,  those  portions  upon  which 
the  metal  will  deposit  should  not  be  touched  with  the  fingers. 

4.  Under  the  conditions  named  in  the  procedure,  the  cop- 
per mny  be  deposited  in  satisfactory  condition  by  a  current 
from  three  cells  of  a  *'  gravity  battery,"  in  series,  or  by  the 
current  which  passes  through  three  lo-candle  Edison  lamps, 
in   series,   on  a    no- volt  circuit.     The   deposition  is  usually 
complete  after  twelve  to  fifteen  hours. 

It  is  not  well  to  leave  the  solution  in  the  circuit  for  an  ex- 
cessive length  of  time,  since,  after  the  deposition  of  the  cop- 
per, the  nitric  acid  is  slowly  reduced  to  ammonia  by  the 
nascent  hydrogen,  with  the  possible  production  of  an  alkaline 
solution  from  which  zinc  may  be  thrown  down. 

5.  The   electrodes  should   be   washed  as  far  as   possible 
before   the   current  is  broken,   to  prevent  re-solution  of  the 
copper.     If  several  solutions  are  connected  in  the  same  cir- 
cuit, some  provision  must  be  made  by  which  the  breaking  of 
the  current  shall  be  avoided  when  the  electrodes  in  any  one 
solution  are  removed.     This  can  be   easily  accomplished  by 
springing  a  piece  of  brass  between  the  binding  posts,  before 
removing  the  electrodes.     The  current  at  first  passes  through 
both  spring  and  solution,  and  when  the  current  is  broken  in  the 
latter,  it  passes  through  the  spring  instead,  and  no  copper  is 
•dissolved  from  the  electrodes  in  other  solutions. 

6.  The  electrode  is  washed  with  alcohol  to  promote  rapid- 
ity of  drying.     The  copper  should  not  remain  in  the  hot  closet 
a  moment  longer  than  is  necessary,  as  it  tends  to  oxidize  at 
the  higher  temperature. 

7.  A  dark  deposit  on  the  anode  indicates  a  precipitation 


DETERMINATION  OF  ZINC.  55 

of  lead  as  peroxide.  Such  deposition  is  not  infrequent,  as 
the  lead  sulphate  is  not  absolutely  insoluble  in  the  acid  liquid 
from  which  it  separates.  This  electrode  may  be  weighed 
with  the  precipitate,  then  cleansed  and  again  weighed,  and 
the  amount  of  lead  calculated  from  the  weight  of  the  perox- 
*  ide  added  to  that  found  as  sulphate  in  the  corresponding 
solution. 

For  a  further  statement  of  the  properties  of  ccpper  and  of 
lead  peroxide,  the  student  is  referred  to  Frtsenius*  Quantita- 
tive Analysis,  under  "  Forms.' 

DETERMINATION    OF    ZINC. 

Procedure.  —  Transfer  the  solution,  from  which  the  copper 
has  been  removed,  to  large  porcelain  casseroles  and  dilute 
to  300  cc.  Cover  these  casseroles  and  add  sodium  carbonate 
solution  to  distinctly  alkaline  reaction ;  boil  until  no  odor 
of  ammonia  can  be  detected  from  the  hot  solution,  adding 
water  to  replace  loss  by  evaporation.  Again  add  i  cc.  or 
2  cc.  of  the  sodium  carbonate  solution  and  boil  to  insure 
complete  removal  of  the  ammonia.  Filter  off  the  basic  zinc 
carbonate,  using  two  niters  for  each  determination,  if  the 
precipitate  is  too  bulky  to  be  properly  held  by  one ;  wash 
the  precipitates  three  times  by  decantation  with  hot  water, 
and  finally  on  the  filter  until  freed  from  carbonate. 

Test  the  filtrate  and  washings  for  complete  precipitation 
by  the  addition  of  a  few  drops  of  ammonium  sulphide,  and 
allow  it  to  stand  ;  if  zinc  sulphide  separates,  collect  it  on 
a  filter,  wash  with  water  containing  a  very  little  ammonium 
sulphide,  ignite  in  a  porcelain  crucible,  and  weigh  as  oxide. 

Dry  the  precipitated  basic  carbonate  in  the  drying  closet, 
and  separate  the  carbonate  as  completely  as  possible  from 
the  filter,  as  described  on  page  24.  Clean,  heat,  and  weigh 
a  No.  7  porcelain  crucible,  burn  the  filter  on  a  wire  above  it, 
allow  the  ash  to  fall  into  the  crucible,  and  ignite  until  the 
carbon  is  all  destroyed.  Transfer  the  main  portion  of  the 
precipitate  to  the  crucible  and  ignite  at  a  red  heat,  until 
the  weight  is  constant.  From  the  weight  of  zinc  oxide, 
calculate  the  weight  of  zinc  and  the  percentage  of  the  latter 
in  the  brass,  remembering  that  but  one  tenth  of  the  solution 
was  used  for  this  determination. 


5 6  GRAVIMETRIC  ANALYSIS. 

Notes.  —  i.  Porcelain  or  platinum  vessels  are  required 
when  strongly  alkaline  solutions  are  to  be  boiled,  as  in  this 
determination,  to  prevent  the  introduction  of  silica  from  the 
glass  into  the  solution. 

2.  Sodium  carbonate  throws  down  a  basic  zinc  carbonate, 
the  exact  composition  of  which  probably  varies  with  varying 
conditions.     This  precipitate  is  partially  soluble  in  the  pres- 
ence of  either  ammonia  or  ammonium  salts,  and  in  the  pres- 
ence of  carbon  dioxide.     Both  of  these  may  be  removed  by 
continued  boiling  with  an  excess  of  sodium  carbonate. 

The  precipitate  is  frequently  very  voluminous,  though  it 
varies  widely  in  this  respect  under  apparently  similar  condi- 
tions. The  student  is  therefore  not  to  suppose  that  errors 
have  been  made  if  one  precipitate  should  appear  much  larger 
than  the  other. 

3.  The  filtrate  should  always  be   tested  for  zinc  by  the 
addition  of  ammonium  sulphide ;  but  when  the  boiling  before 
filtration   is  sufficiently  prolonged,  no  weighable  quantity  of 
zinc  should  be  found  in  the  filtrate.      The  zinc  sulphide,  if 
any  appears,  must  be  washed  with  water  containing  ammo- 
nium sulphide  to  prevent  oxidation. 

4.  The  basic  zinc  carbonate  loses  water  and  carbon  diox- 
ide on  ignition,  leaving  the  zinc  oxide.     This  oxide  may  be 
reduced  to  metallic  zinc  by  contact  with  burning  carbon,  and 
the  zinc  may  volatilize ;  a  separation  of  the  main  portion  of 
the  precipitate  from  the  filter  is,  therefore,  necessary.     This 
separation  should  be  made  as  complete  as  possible,  without 
an  admixture  of  shreds  from  the  filter  paper. 

It  is  well  to  test  the  ignited  zinc  oxide,  after  weighing;  with 
moist  litmus  paper  for  alkaline  reaction.  If  this  is  found,  the 
presence  of  alkali  in  the  precipitate  is  indicated.  It  should 
then  be  boiled  with  water  and  again  ignited,  with  the  same 
precaution  as  before. 

5.  The  zinc  may  also  be  determined  in  the  solution  from 
which  the  copper  has  been  removed,  by  precipitation  as  zinc 
ammonium  phosphate  and  ignition  to  pyrophosphate,  a  method 
preferred  by  many  to  the  carbonate  precipitation.     For  a  de- 
scription of  this  procedure  the  student  is  referred  to  /.  Soc. 
Chem.  Ind.,  15,  866. 

For  a  further  statement  of  the  properties  of  basic  zinc  car- 
bonate and  of  zinc  oxide,  the  student  is  referred  to  Fresenius* 
Quantitative  Analysis,  under  "  Forms." 


LAUMQNTITE  AND  FELDSPAR.  57 

DETERMINATION    OF   SILICA   IN    LAUMONTITE 
AND    IN    FELDSPAR. 

Of  the  natural  and  artificially  prepared  silicates  a  com- 
paratively few  are  completely  decomposed  by  acids,  while 
a  large  number  require  to  be  disintegrated  by  fusion  before 
complete  solution  can  be  accomplished. 

The  mineral  laumontite  is  taken  as  a  type  of  the  decom- 
posable silicates,  and  feldspar  of  those  requiring  fusion  for 
disintegration. 

ANALYSIS    OF    LAUMONTITE. 

Grind  the  mineral  to  a  fine  powder.  Weigh  out  two  por- 
tions of  0.5  to  0.6  gram  each  into  300  cc.  porcelain  casse- 
roles. Pour  over  them  15  cc.  of  water,  and  stir  until  the 
powder  is  evenly  mixed  with  the  water.  Add  25  cc.  of  hydro- 
chloric acid  (1.12  sp.  gr.)  in  small  portions,  and  warm  until 
the  mineral  is  wholly  disintegrated,  as  indicated  by  a  floc- 
culent  residue  and  absence  of  grittiness  under  the  stirring 
rod  (if  any  remains  un dissolved).  Evaporate  the  solution 
to  dryness,  add  15  cc.  of  hydrochloric  acid  (sp.  gr.  1.12),  and 
again  evaporate,  and  heat  the  residue  for  two  hours  at  a 
temperature  approximating  130°  C.  to  dehydrate  the  silicic 
acid.  Moisten  the  residue  with  hydrochloric  acid  (1.12  sp. 
gr.),  warm  gently,  making  sure  that  the  acid  comes  into 
contact  with  the  whole  of  the  residue,  dilute  to  200  cc.,  and 
bring  to  boiling.  Filter  off  the  silica,  wash  four  or  five 
times  with  warm  dilute  hydrochloric  acid  (one  part  acid  [1.12 
sp.  gr.]  to  three  parts  of  water)  and  finally  with  hot  water 
until  free  from  hydrochloric  acid.  Transfer  the  filter  to  a 
platinum  crucible  and  ignite,  as  described  on  page  16,  to 
constant  weight. 

Cautiously  pour  into  the  crucible  about  0.5  cc.  of  hydro- 
fluoric acid.  This  must  be  done  in  a  hood,  with  a  good 
draught,  and  great  care  must  be  taken  that  the  acid  does  not 
come  into  contact  with  the  hands,  as  it  produces  painful 
wounds. 


58  GRAVIMETRIC  ANALYSIS. 

If  the  precipitate  has  dissolved  in  this  quantity  of  acid, 
add  two  drops  of  concentrated  sulphuric  acid,  and  heat  very 
slowly  (always  under  the  hood)  until  all  the  liquid  has  evapo- 
rated, finally  igniting  to  redness.  Cool  in  a  desiccator,  and 
weigh  the  residue.  Deduct  the  weight  of  this  residue  from 
the  previous  apparent  weight  of  silica,  and  from  the  differ- 
ence calculate  the  percentage  of  silica  present. 

Notes. —  i.  Laumontite  is  a  hydrated  silicate  of  aluminum 
and  calcium.  Other  decomposable  silicates  or  slags  may  be 
substituted  for  this  mineral,  as  the  procedure  is  of  general 
application. 

2.  If  strong  acid  were  poured  directly  upon  the  powdered 
silicate  a  partial  separation  of  gelatinous  silicic  acid  would 
occur,  and  this  jelly-like  mass  would  inclose  particles  of  the 
unchanged   mineral  protecting  them  from  the  action   of  the 
acid.     The  water  is  added  to  avoid  this  separation  by  distrib- 
uting the  particles  and  diluting  the  first  portions  of  acid. 

3.  A  flocculent  residue  will  often  remain  after  the  decom- 
position of  the  mineral  is  effected.     This  is  usually  partially 
dehydrated  silicic  acid.     Silicic  acid  is  only  completely  held 
in  solution  by  acids  when  it  is  in  the  fully  hydrated  condition, 
corresponding  to  the  symbol   Si(OH)4.      This  compound  is 
very  unstable,  and  soon  loses  water,  separating  from  solutions 
as  the  loss   increases,   until,   on  protracted  heating  at  tem- 
peratures above  100°  C.,  it  loses  all  its  water,  and  becomes 
practically  insoluble.      The   progress  of   the  dehydration  is 
indicated  by  the  behavior  of  the  solution,  which  as  evapo- 
ration  proceeds  usually  gelatinizes.      On   this  account  it  is 
necessary  to  allow  the  solution  to  evaporate  on  a  steam  bath, 
or  to  stir  it  vigorously,  to  avoid  loss  by  spattering. 

4.  To  obtain   pure  silica,  the   residue,  after  evaporation, 
must  be  thoroughly  extracted  by  warming  with  hydrochloric 
acid,   and  the  solution   freely   diluted  to   prevent,   as  far  as 
possible,  the  inclosure  of  the  residue  in  the  particles  of  silica. 

5.  Aluminum  and  iron  are  likely  to  be  thrown  down  as 
basic  salts  from  hot,  very  dilute  solutions  of  their  chlorides 
If  the  silica  were  washed  only  with  hot  water,  the  solution  of 
these  chlorides  remaining  in  the  filter  after  the  passage  of  the 
original  filtrate,  would  gradually  become  so  dilute  as  to  throw 


LAUMONTITE   AND  FELDSPAR.  59 

down  basic  salts  within  the  pores  of  the  filter,  which  would 
remain  with  the  silica.  To  avoid  this,  an  acid  wash  water  is 
used  until  the  aluminum  and  iron  are  practically  removed. 
The  acid  is  then  removed  by  water. 

6.  The  silica  undergoes   no  change   during   the    ignition 
beyond  the  removal  of  all  traces  of  water. 

7.  Notwithstanding  all  precautions,  the  ignited  precipitate 
of  silica  is  rarely  wholly  pure.     It  is  tested  by  volatilization 
of  the  silica  as  silicon  fluoride  after  solution  in  hydrofluoric 
acid,  and,  if  the  analysis  has  been   properly  conducted,  the 
residue,  after  treatment  with  the  acids  and  ignition,  should 
not  exceed  one  milligram. 

The  acid  produces  ulceration  if  brought  into  contact  with 
the  skin,  and  its  fumes  are  excessively  harmful  if  inhaled. 

8.  The  impurities  are  probably  weighed  with  the  original 
precipitate  as  oxides.      The   addition   of  the   sulphuric   acid 
displaces  the  hydrofluoric   acid,  and  the   resulting  sulphates 
(usually  of  iron  or  aluminum)  are  converted  to  oxides  by  the 
final  ignition. 

It  is  obvious  that  unless  the  acids  used  are  known  to  leave 
no  residue  on  evaporation,  a  quantity  equal  to  that  employed 
in  the  analysis  must  be  evaporated,  and  a  correction  applied 
for  any  residue  found. 

9.  According  to  Gilbert  (Technology   Quarterly,  3,  61)  it 
is  frequently  impracticable  to  render  silicic  acid   completely 
insoluble  by  action  of  heat  alone.     The  silica  which  passes 
into  solution  is  thrown   down   on  the   addition   of  ammonia, 
with  the  iron  and  aluminum,  and  may  be  recovered  by  ignit- 
ing this  precipitate,  grinding  it  carefully  in  an  agate  mortar, 
fusing  it  with  acid  potassium   sulphate,  dissolving  the  fused 
mass  in  water,  and  filtering  off  and  weighing  the  silica. 

The  amount  of  silica  which  fails  of  determination  by  the 
heating  alone,  as  prescribed  above,  is  very  small,  and  is,  in 
this  instance,  lessened  by  the  presence  of  calcium  chloride 
from  the  laumontite  itself.  This  additional  procedure  for  the 
recovery  of  the  soluble  silica  may  be  omitted  unless  extreme 
accuracy,  as  in  certain  mineral  analyses,  is  essential. 

For  a  further  statement  of  the  properties  of  silicic  acid 
and  of  silica  the  student  is  referred  to  Fresmius*  Quantitative 
Analysis,  under  "Forms." 


60  GRAVIMETRIC  ANALYSIS. 

ANALYSIS    OF    FELDSPAR. 

Grind  about  3  grams  of  the  mineral  in  an  agate  mortar 
until  no  grittiness  is  to  be  detected  when  the  mineral  is 
placed  between  the  teeth,  or,  better,  until  it  will  entirely 
pass  through  a  sieve  made  of  fine  silk  bolting  cloth.  The 
sieve  may  be  made  by  placing  a  piece  of  the  bolting  cloth 
over  the  top  of  a  small  beaker  in  which  the  ground  mineral 
is  placed,  holding  the  cloth  in  place  by  means  of  a  rubber 
band  below  the  lip  of  the  beaker.  By  inverting  the  beaker 
over  clean  paper  and  gently  tapping  it,  the  fine  particles  pass 
through  the  sieve,  leaving  the  coarser  particles  within  the 
beaker.  These  must  be  returned  to  the  mortar  and  ground 
until  they  will  pass  through  the  sieve. 

Weigh  out  into  platinum  crucibles  two  portions  of  the 
ground  feldspar  of  about  0.8  gram  each.  Weigh  out  on 
rough  balances  two  portions  of  anhydrous  sodium  carbonate, 
each  amounting  to  about  six  times  the  weight  of  the  feldspar 
taken  for  analysis.  Pour  about  three  fourths  of  the  sodium 
carbonate  into  the  crucible,  place  the  latter  on  a  piece  of 
clean,  glazed  paper,  and  thoroughly  mix  the  substance  and 
the  flux  by  carefully  stirring  for  several  minutes  with  a  dry 
glass  rod.  The  rod  may  be  wiped  off  with  a  small  fragment 
of  filter  paper,  which  may  be  placed  in  the  crucible.  Place 
the  remaining  fourth  of  the  carbonate  on  the  top  of  the 
mixture.  Cover  the  crucible,  heat  it  to  dull  redness  for 
five  minutes,  and  then  increase  the  heat  to  the  full  capacity 
of  a  Bunsen  or  Tirrill  burner  for  twenty  minutes,  or  until 
a  quiet  liquid  fusion  is  obtained.  Finally  heat  the  sides  and 
cover  strongly,  until  any  material  which  may  have  collected 
upon  them  is  also  brought  to  fusion. 

Allow  the  crucible  to  cool,  and  remove  the  fused  mass  as 
directed  on  page  32.  Disintegrate  the  fused  mass  by  plac- 
ing it  in  a  previously  prepared  mixture  of  100  cc.  of  water 
and  50  cc.  of  hydrochloric  acid  (sp.  gr.  1.12),  in  a  covered 
casserole.  Clean  the  crucible  and  lid  by  means  of  a  little 
hydrochloric  acid,  adding  this  acid  to  the  main  solution. 
When  disintegration  is  complete,  evaporate  the  solution  to 


LAUMONTITE  AND  FELDSPAR.  6r 

dryness,   heat   the   residue  to    130°   C.  for  two   hours,   and 
proceed  as  in  the  analysis  of  laumontite. 


i.  If  the  feldspar  is  in  the  massive  or  crystalline 
form,  clean  pieces  should  be  selected,  and  crushed  in  an  iron 
mortar  to  about  half  the  size  of  a  pea,  and  then  transferred  to 
a  steel  mortar,  in  which  they  are  reduced  to  a  coarse  powder 
ready  for  the  agate  mortar.  A  wooden  mallet  should  always 
be  used  to  strike  the  pestle  of  the  steel  mortar,  and  the  blows 
should  not  be  sharp. 

It  is  plain  that  the  grinding  must  continue  until  the  whole 
of  the  portion  of  the  mineral  originally  taken  has  been  ground 
so  that  it  will  pass  the  bolting  cloth,  otherwise  the  sifted  por- 
tion does  not  represent  an  average  sample,  the  softer  ingredi- 
ents, if  foreign  matter  is  present,  being  first  reduced  to  powder. 
For  this  reason  it  is  best  to  start  with  about  the  quantity  of 
the  feldspar  needed.  Any  coarse  particles  remaining  would 
resist  the  action  of  the  flux. 

2.  During  the   fusion   the  feldspar,  which   is  a  strictly  a 
silicate   of   aluminum   and  either  sodium   or   potassium,   but 
usually  contains  some  iron,  calcium,  and  magnesium,  is  de- 
composed by  the  alkaline  flux,  the  sodium  of  the  latter  com- 
bining with  the  silicic  acid  of  the  silicate  with  the  evolution  of 
carbon  dioxide,  while  about  two-thirds  of  the  aluminum  forms 
sodium  aluminate,  and  the  remainder  is  converted  to  a  basic 
carbonate,   or  the   oxide.      The   calcium   and   magnesium,  if 
present,    are   changed   to  carbonates,    and,    with    the    alumi- 
num   oxide,    remain    undissolved    on    treatment   with   water. 
The  heat  is  applied  gently,  to   prevent  a  too  violent  reac- 
tion at  the  start. 

3.  The  silicic  acid   must  be  freed  from   its  combination 
with  a  base  (sodium,  in  this  instance)  before  it  can  be  dehy- 
drated.     The  excess  of  hydrochloric  acid  accomplishes  this 
liberation.     By  disintegrating  the  fused  mass  with  a  consider- 
able volume  of  dilute  acid  the  silicic  acid  is  at  first  held  in 
solution  to  a  large  extent,  and  on  evaporation  and  re-solution 
of  the  bases,  the  silica  is  left  in  a  better  condition  for  filtration 
and  washing.     Immediate  treatment  of  the  fused  mass  with 
strong  acid  is  likely  to  cause  the  silicic  acid  to  separate  at 
once,  and  to  inclose  alkali  salts,  or  alumina. 

4.  A  portion  of  the  fused  mass  is  usually  projected  up- 


62  GRAVIMETRIC  ANALYSIS. 

ward  by  the  escaping  carbon  dioxide  during  the  fusion.  The 
crucible  must  therefore  be  kept  covered  as  much  as  possible, 
and  the  lid  carefully  cleaned. 

5.  A  gritty  residue  remaining  after  the  disintegration   of 
the  fused  mass  by  water  indicates  that  the  substance  has  been 
but  imperfectly  decomposed.     Such  a  residue  should  be  fil- 
tered, washed,  dried,  ignited,  and  again  fused  with  the  alka- 
line flux;  or,  if  the  quantity  of  material  at  hand  will  permit,  it 
is  better  to  reject  the  analysis,  and  to  use  increased  care  in 
grinding  the  mineral  and  in  mixing  it  with  the  flux. 

A  large  residue  remaining  after  the  volatilization  of  the 
silica  also  indicates  imperfect  decomposition  of  the  feldspar 
by  the  fusion. 

6.  Quartz,  and  minerals  containing  very  high  percentages 
of  silica,  may  require  eight  or  ten   parts  by  weight  of  the 
flux  to  insure  a  satisfactory  decomposition. 

7.  The  presence   of   the   large   quantity  of   sodium    salts 
tends  to  increase  the   amount   of   silica  which   fails   to   be- 
come insoluble  during  the  heating  (Gilbert,  loc.  «'/.).      For 
work  of  the  highest  accuracy  the  remaining  silica  should  be 
recovered,  as  outlined  on  page  59,  note  9. 


PART   III. 


VOLUMETRIC     ANALYSIS. 


GENERAL    DISCUSSION. 

IT  has  already  been  pointed  out  in  Part  I,  that  the  meas- 
urement of  the  volume  of  a  solution  required  for  a  definite 
reaction,  takes  the  place  in  Volumetric  Analysis  which  is 
occupied  by  the  weighing  of  the  precipitated  body  in  Gravi- 
metric Analysis. 

It  is  plain  that  the  analytical  balance  is  equally  requisite 
as  a  starting  point  for  both  systems  ;  and  it  will  be  seen  that 
the  processes  of  volumetric  analysis  demand,  beside  an  ac- 
curate balance,  standard  solutions ;  i.e.,  solutions  of  accu- 
rately known  value ;  graduated  instruments  in  which  to 
measure  the  volume  of  such  solutions;  and  finally,  some 
means  which  shall  furnish  an  accurate  indication  of  the  point 
at  which  the  desired  reaction  is  completed.  Those  sub- 
stances which  furnish  such  information  are  called  indicators. 
The  last-named  will  be  treated  of  in  connection  with  the  dif- 
ferent analyses. 

The  process  whereby  a  standard  solution  is  brought  into 
reaction  is  called  titration,  and  the  point  at  which  the  re- 
action is  exactly  completed  is  called  the  end-point.  The 
indicator  should  show  the  end-point  of  the  titration. 

The  processes  of  Volumetric  Analysis  are  easily  classified, 
according  to  their  character,  into  : 

I.  Saturation   Methods ;    such,    for  example,   as  those    of 
acidimetry  and  alkalimetry. 

II.  Oxidation  Processes  ;  as  exemplified  in  the  determina- 
tion of  ferrous  iron,  by  its  oxidation  with  potassium  bichro- 
mate. 

III.  Precipitation    Methods ;    of   which    the    titration   for 
silver  with  potassium  sulphocyanate  solution  is  an  illustration. 


64  VOLUMETRIC  ANALYSIS, 

From  a  somewhat  different  standpoint  the  methods  may 
be  sub-divided  into,  (a)  Direct  Methods,  in  which  the  sub- 
stance sought  is  directly  determined  by  titration  with  a 
standard  solution  to  an  end-point ;  and  (b)  Indirect  Methods^ 
in  which  the  substance  itself  is  not  measured,  but  a  quantity 
of  reagent  known  to  be  an  excess,  with  respect  to  a  specific 
reaction,  is  added,  and  the  unused  excess  determined  by 
titration.  Examples  of  the  latter  class  will  be  pointed  out 
as  they  occur  in  the  procedures. 

Volumetric  processes  are,  as  a  rule,  more  rapid  and  fre- 
quently more  accurate  than  gravimetric  processes  having  the 
same  ends  in  view.  The  number  of  reactions  capable  of 
adaptation  as  volumetric  methods  is,  however,  somewhat 
limited. 

STANDARD    SOLUTIONS. 

The  strength  or  value  of  a  solution  for  a  specific  reaction 
is  determined  by  a  procedure  called  Standardization,  in 
which  the  solution  is  brought  into  reaction  with  a  definite 
weight  of  a  substance  of  known  purity.  For  example,  a 
definite  weight  of  pure  sodium  carbonate  may  be  dissolved 
in  water,  and  the  volume  of  a  solution  of  hydrochloric  acid 
necessary  to  exactly  neutralize  the  carbonate,  accurately  de- 
•termined.  From  these  data  the  strength  or  value  of  the 
acid  is  known.  It  is  then  a  standard  solution. 

Standard  solutions  may  be  made  of  a  purely  empirical 
strength,  dictated  solely  by  convenience  of  manipulation,  or 
the  concentration  may  be  chosen  with  reference  to  a  system 
which  is  applicable  to  all  solutions,  and  based  upon  chemical 
equivalents.  Such  solutions  usually  bear  some  simple  rela- 
tion to  a  Normal  Solution  of  the  specific  reagent ;  i.  e.,  they 
are,  for  example,  deci-normal  or  centi-normal  solutions. 

A  Normal  Solution,  as  defined  by  Mohr  (Titrirmethode, 
p.  56),  contains  in  one  liter  "  one  equivalent  of  the  active  re- 
agent in  grams."  The  "  equivalent  in  grams "  may  be 
defined  as  "  that  quantity  of  the  active  reagent  which  con- 
tains, replaces,  unites  with,  or  in  any  way,  directly  or  indi. 
rectly,  brings  into  reaction  one  gram  of  hydrogen." 

The  application  of  this  general  statement  to  specific  cases 
is  pointed  out  below. 


NORMAL   SOLUTfONS.  65 

A  liter  of  normal  acid  solution  should  contain  such  a 
quantity  of  the  reagent  as  will  furnish  I  gram  of  hydrogen 
replaceable  by  a  base.  Accordingly,  the  normal  solution  of 
hydrochloric  acid  (HC1)  should  contain  36.45  grams  of  the 
gaseous  compound,  since  that  amount  furnishes  the  requisite 
i  gram  of  hydrogen.  On  the  other  hand,  the  normal  solu- 
tion of  sulphuric  acid  (H2SO4)  should  contain  only  49.04 
grams,  one  half  of  its  molecular  weight. 

A  normal  alkali  solution  should  contain  sufficient  alkali  in 
a  liter  to  replace  I  gram  of  hydrogen  in  an  acid.  This  quan- 
tity is  represented  by  the  molecular  weight  in  grams  (40.05) 
of  sodium  hydrate  (NaOH),  while  a  sodium  carbonate  solu- 
tion (Na2CO3)  should  contain  but  one  half  the  molecular 
weight  in  grams  (i.e.,  53.05  grams)  in  a  normal  solution. 

A  normal  solution  of  an  oxidizing  agent  should  contain 
one  equivalent  of  available  oxygen  ;  that  is,  sufficient  oxygen 
to  unite  with  I  gram  of  hydrogen  to  form  water.  The 
amount,  for  example,  of  potassium  bichromate  (K2Cr2O7), 
which  will  furnish  one  equivalent  of  available  oxygen  is 
seen  from  the  following  considerations  :  The  bichromate 
yields  on  reduction,  a  salt  of  potassium,  corresponding  to  the 
oxide  K2O,  and  a  salt  of  chromium,  corresponding  to  the 
oxide  Cr2O3.  The  residual  and  available  oxygen  is  repre- 
sented by  three  atoms  (K2Cr2O7  =  K2O  +  Cr2O3  +  O3). 
Accordingly,  i  gram-molecule  of  the  bichromate  will  fur- 
nish six  equivalents  of  oxygen,  or  enough  to  oxidize  6  grams 
of  hydrogen  to  water,  as  seen  from  the  equation  6H2  -f-  3O2 
=  6H2O.  The  definition,  therefore,  demands  only  one  sixth 
of  the  molecular  weight  (or  49.08  grams)  for  a  normal 
solution. 

A  liter  of  a  normal  solution  of  a  reducing  agent  must 
have  the  same  reducing  power  as  i  gram  of  hydrogen.  For 
example,  a  solution  of  stannous  chloride  must  contain  one 
half  of  its  molecular  weight  in  grams  per  liter  (94.97  grams), 
as  indicated  by  the  equations,  SnCl2  -f-  2FeCls  =  SnCl4  + 
2FeCl2  and  H2  +  2FeCl3  =  2HC1  +  2FeCl2.  One  gram- 
molecule  of  the  stannous  chloride  is  plainly  equivalent  to  2 
grams  of  hydrogen. 

Normal  solutions,  since  they  rest  upon  a  common  founda- 


66  VOLUMETRIC  ANALYSIS. 

tion,  have  the  advantage  of  uniformity.  A  liter  of  a  normal 
solution  of  an  acid  must,  of  necessity,  exactly  neutralize  a 
liter  of  normal  alkali  solution,  and  one  of  an  oxidizing  agent 
exactly  react  with  a  liter  of  a  normal  reducing  solution,  and 
so  on.  It  must  at  the  same  time  be  remembered  that  the 
same  substance  may  have  different  equivalents  when  used 
under  varying  conditions  ;  as,  for  example,  potassium  perman- 
ganate, two  molecules  of  which  yield  three  atoms  of  availa- 
ble oxygen  (equivalent  to  six  hydrogen  atoms)  in  neutral 
solution,  and  five  atoms  of  oxygen  (equivalent  to  ten  atoms 
of  hydrogen)  when  used  in  acid  solution.  These  facts  must 
be  considered  when  reference  is  made  to  a  normal  solution 
of  that  reagent,  and  the  statement  must  specify  the  condi- 
tions of  use. 

Beside  the  advantage  of  uniformity,  the  use  of  normal 
solutions  simplifies  the  calculations  of  the  results  of  analyses. 
This  is  particularly  true  if,  in  connection  with  the  normal 
solution,  the  weight  of  substance  for  analysis  be  chosen 
with  reference  to  the  molecular  weight  of  the  constituent 
to  be  determined.  For  illustrations  of  this,  consult  Part  V, 
page  134. 

The  preparation  of  an  exactly  normal,  half  normal,  or  deci- 
normal  solution  requires  considerable  time  and  care,  as  noted 
on  page  78,  and  is  usually  carried  out  when  a  large  number 
of  analyses  are  to  be  made,  or  when  the  analyst  has  some 
other  specific  purpose  in  view.  It  is,  however,  a  compara- 
tively easy  matter  to  prepare  standard  solutions  which  differ 
but  slightly  from  the  normal  or  half  normal  solutions,  and 
these  have  the  advantage  of  practical  equality.  That  is,  two 
approximately  half  normal  solutions  are  more  convenient  to 
work  with,  than  two  which  are  widely  different  in  strength. 
It  is,  however,  true,  that  whatever  advantage  pertains  to  the 
use  of  normal  solutions  as  regards  simplicity  of  calculations 
is,  to  a  considerable  extent,  lost  when  using  these  approxi- 
mate solutions. 

GRADUATED    INSTRUMENTS. 

A  burette  consists  of  a  glass  tube  which  is  made  as  uni- 
formly cylindrical  as  possible,  and  of  such  a  bore  that  the 


CALIBRA  TION.  67 

divisions  which  are  etched  upon  its  surface  shall  correspond 
to  actual  contents  as  far  as  is  practicable. 

The  tube  is  contracted  at  one  extremity,  and  terminates 
in  either  a  glass  stopcock  and  delivery  tube  (a  Giessler  bu- 
rette), or  in  such  a  manner  that  a  piece  of  rubber  tubing 
may  be  firmly  attached,  connecting  a  delivery  tube  of  glass. 
The  rubber  tubing  is  closed  by  means  of  a  pinchcock  or  by 
a  glass  bead  (Mohr  burette). 

The  graduations  are  usually  numbered  in  cubic  centime- 
ters, and  the  latter  are  subdivided  into  tenths. 

A  pipette  may  consist  of  a  narrow  tube,  in  the  middle  of 
which  is  blown  a  bulb  of  a  capacity  a  little  less  than  that 
which  it  is  desired  to  measure  by  the  pipette  ;  or  it  may  be  a 
miniature  burette,  without  the  stopcock  or  rubber  tip  at  the 
lower  extremity.  In  either  case,  the  flow  of  liquid  is  regu- 
lated by  the  pressure  of  the  finger  on  the  top,  which  pre- 
vents the  admission  of  the  air. 

Graduated,  or  measuring  flasks  are  similar  to  the  ordinary 
flat-bottomed  flasks,  but  are  provided  with  long,  narrow 
necks  in  order  that  slight  variations  in  the  position  of  the 
meniscus  with  respect  to  the  graduation,  shall  represent  a 
minimum  volume  of  liquid.  The  flasks  must  be  of  such  a 
capacity  that,  when  filled  with  the  specified  volume,  the 
liquid  rises  well  into  the  neck. 

CALIBRATION    OF     INSTRUMENTS. 

If  accuracy  of  results  is  to  be  attained,  the  correctness  of 
all  measuring  instruments  must  be  tested.  None  of  the  ap- 
paratus offered  for  sale  can  be  implicitly  relied  upon,  unless 
it  be  those  more  expensive  instruments  which  are  accompa- 
nied by  a  certificate  from  the  Physikalische  Reichsanstalt  in 
Berlin,  or  other  equally  authentic  source. 

The  bore  of  burettes  may  readily  vary,  and  as  the  gradua- 
tions must  be  applied  without  regard  to  such  variations  of 
bore,  local  errors  are  the  result.  The  same  consideration  ap- 
plies to  pipettes,  while  even  the  graduations  upon  flasks  are 
often  incorrect  for  the  temperatures  given.  It  is  the  custom 
in  most  laboratories  to  purchase  the  flasks  ungraduated,  and 
to  graduate  them  for  the  standard  in  use. 


68  VOL UME TRIG  ANAL  YSSS. 

The  process  of  testing  these  instruments  is  called  Calibra- 
tion. It  is  usually  accomplished  by  comparing  the  actual 
weight  of  water  contained  in  the  instrument  with  its  appar- 
ent volume. 

There  is,  unfortunately,  no  uniform  standard  of  volume 
which  has  been  adopted  for  general  use  in  all  laboratories. 
It  has  been  variously  proposed  to  consider  the  volume  of 
1000  grams  of  water  at  4°,  15.5°,  16°,  17.5°,  and  even  20°  C, 
as  a  liter,  for  practical  purposes,  and  to  consider  the  cubic 
centimeter  to  be  one  thousandth  of  that  volume. 

The  true  liter  is  the  volume  of  1000  grams  of  water  at 
4°  C. ;  but  this  is  obviously  a  lower  temperature  than  that 
found  in  our  laboratories,  and  involves  the  constant  use  of 
corrections,  if  taken  as  the  standard.  Mohr  in  his  Titrir- 
methode  adopts  17.5°  C.  as  a  practically  useful  temperature, 
and  the  volume  of  1000  grams  of  water  at  17.5°  C.  is  known 
as  the  "Mohr  liter."  This  temperature,  or  even  15°  C., 
does  not  differ  greatly  from  the  average  temperature  of 
the  laboratory,  and,  except  when  the  highest  accuracy  is 
required,  no  corrections  need  be  applied.  It  is  always  well, 
however,  to  note  the  temperature  at  the  time  of  a  titration, 
and  to  make  proper  corrections  for  unusual  fluctuations. 

When  apparatus  for  graduation  is  purchased  unmarked, 
and  is  calibrated  in  the  laboratory  where  it  is  to  be  used,  any 
temperature  may  be  selected  which  approaches  the  average 
for  that  laboratory,  but  all  instruments  should,  for  conven- 
ience, be  calibrated  at  that  temperature,  whether  flask,  bu- 
rette, or  pipette,  since,  if  a  50  cc.  measuring  flask  is  to  be 
capable  of  measuring  exactly  one  tenth  of  the  liquid  con- 
tained in  a  500  cc.  measuring  flask,  both  must  be  correct  at 
the  same  temperature.  The  spot  selected  for  volumetric 
work  should  be  subject  to  a  minimum  of  temperature  change. 

CALIBRATION    OF    BURETTES. 

Each  student  should  calibrate  at  least  one  burette  as 
follows : 

Procedure.  —  Clean  the  burette  thoroughly  by  pouring 
into  it  a  warm  solution  of  chromic  acid  in  concentrated 


CA LIBRA  TION  OF  B  URE  TTES.  69 

sulphuric  acid.  Stopper  the  burette  and  bring  the  acid  in 
contact  with  its  entire  length  by  shaking.  Pour  the  acid 
back  into  its  receptacle,  and  wash  out  the  burette  thor- 
oughly with  water.  Unless  the  water  runs  from  the  bu- 
rette without  leaving  drops  upon  the  sides,  the  process  must 
be  repeated.  When  clean,  fill  the  burette  with  distilled  water, 
allow  it  to  run  out  through  the  stopcock,  or  rubber  tip,  until 
convinced  that  no  air  bubbles  are  inclosed.  Fill  the  burette 
to  the  zero  mark  and  draw  off  the  liquid  until  the  meniscus 
is  just  below  the  zero  mark.  To  take  the  exact  reading, 
wrap  around  the  burette  a  piece  of  colored  paper,  with  its 
straight,  smooth  edges  held  evenly  together  (color  turned 
inside),  and  held  two  small  divisions  below  the  meniscus. 
Move  the  eye  so  that  the  edge  of  the  paper  at  the  back  of 
the  burette  is  just  hidden  by  that  in  front,  and  note  the 
position  of  the  lowest  point  of  the  meniscus  of  the  water. 
Estimate  the  tenths  of  the  small  divisions,  corresponding 
to  hundredths  of  a  cubic  centimeter,  and  record  the  reading 
in  the  notebook. 

Weigh  a  50  cc.  flat-bottomed  flask  (of  thin  glass),  which 
must  be  dry  on  the  outside,  to  the  nearest  centigram.  Re- 
cord the  weight  in  the  notebook.  Place  the  flask  under  the 
burette,  and  draw  out  into  it  about  10  cc.  of  water,  removing 
any  drop  on  the  tip  by  touching  it  against  the  inside  of  the 
neck  of  the  flask.  Do  not  attempt  to  stop  exactly  at  the  10 
cc.  mark,  but  do  not  vary  more  than  o.  I  cc.  from  it.  Note 
the  time,  and  at  the  expiration  of  three  minutes  (or  longer), 
take  the  reading  upon  the  burette  accurately,  and  record  it  in 
the  notebook.  Meanwhile,  weigh  the  flask  and  water  to 
centigrams  and  record  its  weight.  Draw  off  the  liquid  from 
10  cc.  to  about  20  cc.  into  the  same  flask  without  emptying 
it ;  weigh,  and  at  the  expiration  of  three  minutes  take  the 
reading,  and  so  on  throughout  the  length  of  the  burette. 
When  it  is  completed,  re-fill  the  burette  and  check  the  first 
calibration. 

The  differences  in  readings  represent  the  apparent  vol- 
ume, the  differences  in  weights,  the  true  volumes.  For 
example,  if  an  apparent  volume  of  10.05  was  found  to  weigh 
10.03  grams,  it  may  be  assumed  with  sufficient  accuracy  that 


7 o  VOL  UME  TRIG  ANAL  YSIS. 

the  error  in  that  10  cc.  amounts  to  0.02  cc.,  or  0.002  for  each 
cubic  centimeter.  The  records  may  conveniently  be  made 
in  the  notebook  under  these  headings  : 

Readings  ;    Differences  ;    Weights  ;    Differences  ;    Calcu- 
lated Corrections. 

Notes.—  i.  The  inner  surface  of  the  burette  must  be  abso- 
lutely clean,  if  the  liquid  is  to  run  off  freely.  Chromic  acid 
in  sulphuric  acid  is  usually  found  to  be  the  best  cleansing 
agent,  but  the  mixture  must  be  warm  and  concentrated.  This 
solution  can  be  prepared  by  adding  to  concentrated  commer- 
cial sulphuric  acid  a  few  crystals  of  potassium  bichromate,  and 
i  cc.  of  water.  Warm  the  mixture  gently  and  pour  off  the 
solution.  It  is  convenient  to  have  such  a  solution  ready  at 
hand,  as  burettes  frequently  need  cleaning.  The  rubber  tip 
should  be  removed  before  the  cleansing  agent  is  added. 

2.  It  is  always  necessary  to    insure    the    absence   of   air 
bubbles  in   the   tips   by  running  the   liquid  rapidly  through 
them.     These  bubbles  may  otherwise  escape  during  titration, 
and  vitiate  results. 

3.  To  obtain  an  accurate  reading,  the  eye  must  be  on  a 
level  with  the  meniscus.     This  may  be  attained  by  the  use  of 
a  paper,  or  by  using  a  float.     The  latter  is  useful,  provided  it 
moves  freely  in  the  burette,  but  care  must  be  taken  that  such 
is  the  case,  otherwise  a  float  is  worse  than  useless. 

4.  The  eye  soon  becomes   accustomed  to  estimating  the 
tenths  of  -the  divisions.     If  the  paper  is  held  as  directed,  two 
divisions  below  the  meniscus,  one  whole  division  is  visible  to 
correct  the  judgment.     It  is  not  well  to  attempt  to  bring  the 
meniscus  exactly  to  a  division   mark  on  the  burette.     Such 
readings  are  usually  less  accurate  than   those    in  which  the 
hundredths  are  estimated. 

5.  It  is  obvious  that  it  would  be  useless  to  weigh  the  water 
with  an  accuracy  greater  than  that  of  the  readings  taken  on 
the  burette.     The  latter  cannot  exceed  o.oi  cc.  in  accuracy, 
which  corresponds  to  o.oi  gram. 

The  student  should  clearly  understand  that  all  other  weigh- 
ings except  those  for  calibration,  should  be  made  accurately  to 
o.oooi  gram. 

6.  A  small  quantity  of  liquid  adheres  to  the  side  of  even 
a  clean  burette.     This  slowly  unites  with  the  main  body  of 


C A  LIBRA  TION  OF  FLASKS.  7 1 

liquid,  but  requires  an  appreciable  time.  Three  minutes  is  a 
sufficient  interval,  but  not  too  long,  and  should  be  adopted  in 
every  instance  throughout  the  whole  volumetric  practice, 
before  final  readings  are  recorded. 

7.  Should  the  error  discovered  in  any  interval  of  10  cc. 
on  the  burette  exceed  o.io  cc.,  it  is  advisable  to  weigh  smaller 
portions  (even  i  cc.),  to  locate  the  position  of  the  variation  of 
bore  in  the  tube,  rather  than  to  distribute  the  correction  uni- 
formly over  the  corresponding  10  cc.     The  latter  is  the  usual 
course  for  small  corrections,  and  it  is  convenient  to  calculate 
the  correction  corresponding  to  each  cubic  centimeter  and 
to  record  it  in  the  form  of  a  table  or  calibration  card,  or  to 
plot  a  curve  representing  the  values. 

8.  Burettes    may  also   be   calibrated  by  drawing  off   the 
liquid  in  successive  portions  through  a  5  cc.  pipette  which  has 
been  accurately  calibrated,  as  a  substitute  for  weighing.     If 
many  burettes  are  to  be  examined  this  is  a  more  rapid  method. 

9.  Pipettes  are  calibrated  in  the  same  general  way  as  bu- 
rettes.    They  must  be  cleaned,  and  are  then  filled  with  water, 
and  the  latter  is  drawn  off  and  weighed.     A  definite  interval 
must  be  allowed  for  draining,  and  a  definite  practice  adopted 
as  regards  the  removal  of  the  liquid  which  collects  at  the  end 
of  the  tube,  if  the  pipette  be.  designed  to  deliver  a  specific 
volume  when  emptied.     This  liquid  may,  at  the  end  of  a  defi- 
nite interval,  be  removed  either  by  touching  the  side  of  the 
vessel  or  by  gently  blowing  out  the  last  drops.     Either  prac- 
tice must  be  uniformly  adhered  to. 

CALIBRATION    OF    FLASKS. 

Procedure.  —  Clean  the  flask  and  dry  it  carefully  outside 
and  inside.  Tare  it  accurately,  and  place  on  the  opposite 
balance-pan  the  number  of  grams  corresponding  to  the  vol- 
ume desired  ;  pour  water  into  the  flask  until  the  weight  of 
the  latter  counterbalances  the  weight  on  the  pan.  Remove 
the  flask  from  the  balance,  stopper  it,  place  it  in  a  bath  at 
the  desired  temperature,  say,  17.5°  C.,  and  after  an  hour, 
mark  on  the  neck  with  a  diamond,  the  location-  of  the  lowest 
point  of  the  meniscus. 

Notes.  —  i.      The     allowable    error    in    counterbalancing 
the  water  and  weights  varies  with  the  volume  of  the  flask. 


7 2  VOL  UME  TRIC  ANAL  YSIS. 

It   should    not  exceed   one  ten-thousandth  of   the   weight  of 
water. 

2.  Other  methods  are  used  which  involve  the  use  of  cali- 
brated apparatus,  from   which  the   desired   volume   of  water 
may  be  run  into  the  dry  flask,  and   the   graduation   marked 
directly  upon  it.     For  a  description  of  one  of  these,  the  stu- 
dent is  referred  to  the  Am.   Chem.  J.,  16,  479. 

3.  Flasks  may  be  graduated  either  for  "  contents  "  or  for 
"delivery."     In    the  former  case  they  contain   the  specified 
volume  when  filled  to  the  graduation  ;  in  the  latter  case  the 
flask  will  deliver  the  specified  volume,  if  allowed  to  drain  for  a 
definite  time.     By  placing  two  marks  upon  the  flask  it  may  be 
graduated  for  both  contents  and  delivery. 

To  calibrate  a  flask  for  delivery,  it  should  be  filled  with 
water,  then  emptied  and  allowed  to  drain  for  a  definite  inter- 
val (three  minutes).  It  is  then  tared,  the  requisite  weights 
are  placed  upon  the  balance  pan,  and  water  added  to  counter- 
balance these.  It  is  then  placed  in  a  bath  at  the  required 
temperature  and,  after  an  hour,  marked. 

Flasks  thus  calibrated  will  deliver  a  definite  volume  of  a 
solution  without  being  washed  out.  It  is,  however,  a  more 
general  custom  to  graduate  flasks  for  contents. 

GENERAL    DIRECTIONS. 

It  is  essential  to  the  success  of  analysis  that  uniformity  of 
practice  shall  prevail  throughout  all  volumetric  work,  with 
respect  to  those  matters  which  can  influence  the  accuracy  of 
measurement  of  liquids.  Whatever  conditions  are  imposed, 
for  example,  during  the  calibration  of  a  burette,  pipette,  or 
flask  (notably  the  time  allowed  for  draining),  must  also  pre- 
vail whenever  the  flask  or  burette  is  used. 

The  student  should  be  constantly  watchful  to  insure  par- 
allel conditions  during  both  standardization  and  analysis,  with 
respect  to  the  final  volume  of  liquid  in  which  a  titration 
takes  place.  The  standard  of  the  solution  is  only  accurate 
under  the  conditions  which  prevailed  when  it  was  determined. 

It  is  plain  that  the  standard  solutions  must  be  scrupu- 
lously protected  from  concentration  or  dilution,  after  their 
value  has  been  established.  Accordingly,  great  care  must 
be  taken  to  thoroughly  rinse  out  all  burettes,  flasks,  etc., 


GENERAL  DIRECTIONS.  73 

with  the  solutions  which  they  are  to  contain,  in  order  to 
remove  all  traces  of  water,  or  other  liquid  which  could  act 
as  a  diluent.  It  is  best  to  wash  out  a  burette  at  least  three 
times  with  small  portions  of  a  solution,  allowing  them  to 
run  out  through  the  tip,  before  assuming  that  it  is  in  a  con- 
dition to  be  filled  and  used.  It  is,  of  course,  possible  to  dry 
the  measuring  instruments  in  a  hot  closet,  but  this  is  tedious 
and  unnecessary. 

To  the  same  end,  all  solutions  should  be  kept  stoppered, 
and  away  from  direct  sunlight  or  heat.  The  bottles  should 
be  shaken  before  use,  to  collect  any  liquid  which  may  have 
distilled  from  the  solution  and  condensed  on  the  sides. 

Care  should  be  taken  when  selecting  a  spot  for  volumetric 
work,  that  no  source  of  heat  is  sufficiently  near  to  raise  the 
temperature  of  the  solutions.  The  temperature  should 
always  be  as  near  17.5°  C.  as  is  practicable. 

Much  time  may  be  saved  by  estimating  the  approximate 
volume  of  a  standard  solution  which  will  be  required  for  a 
titration  (if  the  data  are  obtainable),  before  beginning  the 
operation.  It  is  then  possible  to  run  in  rapidly  approxi- 
mately the  required  amount,  after  which  it  is  only  necessary 
to  determine  the  end-point  with  accuracy.  In  such  cases, 
however,  the  knowledge  of  the  amount  probably  to  be  re- 
quired should  never  be  allowed  to  influence  the  judgment 
regarding  the  end-point. 


74  VOLUMETRIC  ANAL  YSIS. 


I.    SATURATION    METHODS. 


ALKALIMETRY    AND    ACIDIMETRY. 


GENERAL     DISCUSSION. 

STANDARD  solutions  of  acid  and  alkali  are  required  for 
these  processes,  together  with  such  indicators  as  will  accu- 
rately designate  the  point  of  saturation. 

Standard  Acid  Solutions  may  properly  be  prepared  from 
either  hydrochloric,  sulphuric,  or  oxalic  acids.  Hydrochloric 
acid  has  the  advantage  of  forming  soluble  compounds 
with  the  alkaline  earths,  but  its  solutions  cannot  be  boiled 
without  loss  of  strength;  sulphuric  acid  solutions  may  be 
boiled  without  loss,  but  the  acid  forms  insoluble  sulphates  of 
three  of  the  alkaline  earths ;  oxalic  acid  can  be  accurately 
weighed  for  the  preparation  of  solutions,  and  its  solutions 
may  be  boiled  without  loss,  but  it  also  forms  insoluble  oxal- 
ates  with  three  of  the  alkaline  earths,  and  cannot  be  used 
with  certain  of  the  indicators. 

Standard  Alkali  Solutions  may  be  prepared  from  sodium 
or  potassium  hydroxide,  sodium  carbonate,  barium  hydrox- 
ide, or  ammonia.  Of  sodium  and  potassium  hydroxide,  it 
may  be  said  that  they  can  be  used  with  all  indicators,  and 
their  solutions  may  be  boiled,  but  they  absorb  carbon  dioxide 
readily,  and  attack  the  glass  of  bottles  ;  sodium  carbonate 
may  be  weighed  directly,  if  its  purity  is  assured,  but  the 
presence  of  the  carbonic  acid  of  the  carbonate  is  a  disad- 
vantage with  many  indicators  ;  barium  hydroxide,  solutions 
may  be  prepared  which  are  entirely  free  from  carbon  diox- 
ide, and  such  solutions  immediately  show  by  precipitation 
any  contamination  from  absorption,  but  the  hydroxide  is  not 
freely  soluble  in  water ;  ammonia  yields  a  clean  solution,  and 
does  not  absorb  carbon  dioxide  as  readily  as  the  caustic  alka- 
lies, but  its  solutions  cannot  be  boiled,  nor  can  they  be  used 
with  all  indicators. 


INDICA  TORS. 


75 


Half-normal  (£)  or  deci-normal  (T^)  solutions  are  employed 
in  most  analyses  (except  in  the  case  of  the  less  soluble  ba- 
rium hydroxide).  Solutions  of  the  latter  strength  are  con- 
venient, when  small  percentages  of  acid  or  alkali  are  to  be 
determined. 

INDICATORS. 

An  indicator,  to  be  of  service  in  acidimetric  processes, 
must  be  a  substance  of  basic  or  acid  character,  which,  like 
litmus,  will  show  by  a  change  of  color,  the  presence  of  the 
slightest  excess  of  free  acid  or  alkali.  The  number  of  or- 
ganic bodies  which  have  been  proposed  as  indicators  is  large, 
but  of  these  a  few  only  have  come  into  general  use.  The 
most  important  among  the  latter  are  presented  in  the  table 
below,  with  their  characteristics  : 


•5 

JS 

.H  3 

•a  3 

. 

_0 

1 

1 

J"s 

|1 

i 

1 

a 

e 

ra  — 

32 

«  aj 

0 

Indicator. 

0 

0     . 

_c  *"* 

-5 

"1 

Sjf 

"i  s 

.r-    "- 

|-s  . 

^1 

i| 

3* 

5* 

"  ci-2 

j|S-2 

D  s 

2* 

Litmus. 

Red. 

Blue 

Unreliable. 

Reliable 

Reliable 

Reliable 

Methyl  Orange. 

Pink. 

Yellow. 

Reliable. 

Unreliable. 

Reliable. 

Unreliable 

Phenolphthalein 

Colorless. 

Pink 

Unreliable. 

Reliable. 

Unreliable, 

Reliable. 

Lacmoid 

Purple  red. 

Blue 

Unreliable. 

Reliable. 

Reliable 

Unreliable.  (?) 

Cochineal. 

Purple  red. 

Blue 

Reliable. 

Reliable 

Reliable 

Unreliable 

Rosolic  acid 

Yellow 

Pink 

Unreliable. 

Reliable. 

Unreliable 

Unreliable  » 

Alizarine 

Yellow. 

Red. 

Unreliable. 

Reliable. 

Reliable. 

Reliable. 

1  Reliable  with  oxalic  acid. 

Litmus  solution,  for  use  as  an  indicator,  must  be  specially 
prepared.  For  the  details  of  its  preparation  and  preserva- 
tion, the  student  is  referred  to  Suttoris  Volumetric  Analysis, 
under  "Indicators."  A  general  discussion  of  the  subject  of 
indicators  will  also  be  found  there. 

Methyl  Orange  (also  known  as  Orange  No.  Ill)  is  not 
affected  by  carbonic  acid  or  sulphuretted  hydrogen,  in  cold 
solution,  and  is,  therefore,  the  most  convenient  indicator  for 
use  in  the  presence  of  carbonates  or  sulphides,  unless 


7 6  VOL UME TRIG  ANAL  YSIS. 

organic  acids  are  also  present.  The  customary  strength  for 
the  indicator  solution  is  I  gram  per  liter,  which  may  be  dis- 
solved in  a  few  cubic  centimeters  of  alcohol,  and  the  solution 
diluted. 

Methyl  Orange  can  be  used  in  the  presence  of  borates. 

Phenolphthalein   is   of    special  value  for  the   titration   of 
organic  acids,  provided  ammonia  or  ammonium  salts  are  not 
present.     In  the  presence  of  carbonates  it  can  only  be  used 
in  hot  solution.     The  indicator  solution  contains   I  gram  in* 
100  grams  of  alcohol. 

Lacmoid  solutions  should  contain  I  gram  in  200  cc.  of 
alcohol.  It  shows  a  neutral  reaction  with  the  salts  of  some 
of  the  heavy  metals,  thus  allowing  the  titration  of  free  acid 
in  their  presence. 

Cochineal  solutions  are  prepared  by  treating  three  parts  of 
cochineal,  with  two  hundred  parts  water  and  fifty  parts  alco- 
hol. This  indicator  is  of  special  value  for  use  with  ammonia. 

Of  the  indicators  mentioned  in  the  table,  methyl  orange 
and  lacmoid  are  most  sensitive  to  alkalies,  and  phenolphtha- 
lein  is  most  readily  changed  by  acids.  It  is  possible  to 
obtain  from  the  same  solution  an  acid  reaction  toward  phe- 
nolphthalein  and  a  neutral  or  slightly  alkaline  reaction 
toward  methyl  orange.  It  is  obvious,  then,  that  it  is  desir- 
able to  employ  the  same  indicator  both  for  standardization 
and  analysis. 

Since  the  indicators  must  require  a  certain  quantity  of 
acid  or  alkali  in  excess,  to  cause  the  change  of  color,  it  is 
also  plain  that  only  the  requisite  quantity  (usually  one  or  two 
drops  of  solution)  should  be  used. 

PREPARATION    OF    HALF-NORMAL    SOLUTIONS    OF    HYDRO- 
CHLORIC   ACID    AND    SODIUM    HYDROXIDE. 

Procedure.  —  Calculate  the  number  of  cubic  centimeters 
of  aqueous  hydrochloric  acid  (sp.  gr.  1.12  or  1.2)  required  to 
furnish  36.456  grams  of  the  gaseous  compound.  (For  this 
purpose  consult  the  table  on  p.  146).  Measure  out  a  vol- 
ume of  acid  about  10  per  cent,  in  excess  of  the  calculated 
quantity  into  a  clean  2-liter  bottle,  and  dilute  with  distilled 


PREPARATION   OF  SOLUTIONS.  77 

water  to  an  approximate  volume  of  2000  cc.  Shake  the 
solution  thoroughly  for  at  least  a  minute,  to  insure  uni- 
formity. 

Weigh  out,  upon  the  laboratory  balances,  about  46  grams 
of  sodium  hydroxide.  Dissolve  the  hydroxide  in  water,  and 
dilute  to  2000  cc.  Shake  this  solution  also  for  a  minute. 

Select  two  clean  burettes,  and  fill  them  with  the  solutions, 
after  rinsing  them  out  three  times  with  10  cc.  of  the  solution, 
which  should  be  allowed  to  run  out  through  the  tip  to  insure 
the  displacement  of  all  water  from  that  part  of  the  burette. 
When  the  burettes  are  ready  for  use,  and  all  air  bubbles  dis- 
placed from  the  tip,  note  the  exact  position  of  the  liquid  in 
each,  and  record  the  readings  in  the  notebook.  Run  out 
from  the  burette  about  40  cc.  of  the  acid  into  a  beaker,  and 
add  two  drops  of  a  solution  of  methyl  orange ;  dilute  the 
acid  to  about  100  cc.,  and  run  out  alkali  solution  from  the 
other  burette  until  the  pink  has  given  place  to  a  yellow. 
Wash  down  the  sides  of  the  beaker  with  a  little  distilled 
water,  if  the  solution  has  spattered  upon  them,  return  the 
beaker  to  the  acid  burette,  and  again  restore  the  pink ;  con- 
tinue these  alternations  until  the  point  is  accurately  fixed  at 
which  a  single  small  drop  of  either  solution  serves  to  pro- 
duce a  distinct  change  of  color.  It  is  usually  more  satisfac- 
tory to  select  as  an  end-point  the  appearance  of  the  faintest 
pink  which  can  be  recognized.  If  the  titration  has  occupied 
more  than  three  minutes,  the  readings  of  the  burettes  may 
be  immediately  taken  and  recorded  in  the  notebook. 

Re-fill  the  burettes  and  repeat  the  titration.  Correct  the 
burette  readings  as  indicated  by  the  burette  calibrations,  and 
obtain  the  ratio  of  the  sodium  hydroxide  solution  to  that  of 
hydrochloric  acid  by  dividing  the  number  of  cubic  centime- 
ters of  acid  used,  by  the  number  of  cubic  centimeters  of 
alkali  required  for  neutralization.  The  check  results  should 
not  vary  by  more  than  0.2  per  cent,  of  the  total  ratio. 

When  this  ratio  has  been  fully  established,  weigh  out, 
from  a  weighing  tube,  into  No.  4  lipped  beakers,  two  por- 
tions of  about  i  gram  each  of  pure  calcium  carbonate,  not- 
ing the  weights  exactly  (to  o.oooi  gram)  in  the  notebook. 
Cover  the  carbonate  with  25  cc.  of  water  and  add  two  drops 


7 8  •  VOL UME TRIG  ANAL  YSlS. 

of  methyl  orange  solution ;  fill  the  burettes  and  note  initial 
readings ;  run  the  acid  into  the  beaker,  with  cautious  stir- 
ring, until  the  carbonate  has  dissolved,  avoiding  loss  by 
effervescence,  wash  down  the  sides  of  the  beaker,  and  then 
run  in  alkali  until  the  indicator  becomes  yellow.  Finish  the 
titration  as  described  above.  Note  the  readings  on  the  bu- 
rette after  the  proM^interval,  and  record  them  in  the  note- 
book. From  the  -olpi  then  recorded,  it  is  possible  to  deter- 
mine th^yolume  3?"  hydrochloric  acid  neutralized  by  the 
pure  calehim  carbonate,  and  hence  the  weight  of  gaseous 
hydrochloric  acid  in  each  cubic  centimeter.  The  standardi- 
zation must  be  repeated  until  these  values  agree  within  0.2 
per  cent. 

Compare  the  value  just  obtained  with  the  weight  of  the 
acid  which  should  be  present  in  i  cc.  of  -j-  HCl^  and  dilute 
the  solution  accordingly.  (For  example  :  a  solutibn  contain- 
ing .01900  gram  HC1  per  cubic  centimetefvshould  be  diluted 
according  to  the  proportion  .01900  :  .01822  =  x  :  1000; 
i.  e.,  each  1000  cc.  &t  that  solution  should  be  diluted  to 
1042.3  cc.)  Measure^  off  2  liters  accurately  in  a  graduated 
flask,  observing  precautions  mentioned  on  page  72,  and  add 
the  requisite  volume  of  water  from  a  burette.  F>r9m  the 
known  ratio  of  the  two  solutions,  and  the  known  value  of 
the  acid,  calculate  the  requisite  dilution  for  the  alkali  solu- 
tion, and  add  the  necessary  volume  of  water  to  it  also.  De- 
termine the  ratio  between  these  new  solutions  after  shaking 
thoroughly,  and  re-standardize  the  acid  against  calcium  car- 
bonate. The  values  thus  found  should  not  differ,  from  the 
calculated  values  for  half-normal  solutions,  by  more  than  0.2 
per  cent,  in  either  case.  If  a  greater  variation  be  found, 
the  dilution  and  standardization  must  be  repeated,  before 
the  solutions  are  ready  for  use. 

Notes.  —  i.  The  solutions  are  prepared  of  greater  strength 
than  that  finally  desired,  as  they  are  more  readily  diluted 
than  strengthened.  Commercial  sodium  hydroxide  is  usually 
impure,  and  always  contains  more  or  less  carbonate  ;  a  fur- 
ther allowance  is  therefore  made  for  this  factor  by  placing  the 
weight  taken  at  46  grams  for  the  2  liters.  If  the  hydroxide  is 
known  to  be  pure,  a  lesser  amount  (say  42  grams)  will  suffice. 


.   PREPARATION  OF  SOLUTIONS.  79 

2.  Too  much  care  cannot  be  taken  to  insure  perfect  uni- 
formity of  solutions  before  standardization,  and  thoroughness 
in  this  respect  will  often  avoid  much  waste  of  time.     A  solu- 
tion once  thoroughly  mixed  does  not  alter. 

3.  The  liquid  is  diluted  to  100  cc.  during  standardization 
to  make  the  volume  approximately  equal  to  that  which  will 
prevail  during  analysis.     Compare  remarks  on  page  72. 

4.  The  change  from  yellow  to  pink  is  better  defined  than 
the  reverse  change.     The  end-point  should  be  chosen  exactly 
at  the  point  of   change  ;    any  darker  tint  is   unsatisfactory, 
since  it  is  impossible  to  carry  shades  of  color  in  the  memory 
and  to  duplicate  them  from  day  to  day. 

5.  A  variation  of  0.2  per  cent,  of  the  total  value  is  not 
excessive  in  a  beginner's  work.     It  must  not  be  regarded  as  a 
fixed   standard,  however,  or   as  necessarily  of  general  appli- 
cation. 

6._  The  calcium  carbonate  is  covered  with  water  to  lessen 
the  Violence  of  the  action  of  the  acid.  It  is  usually  well  to 
add  a  moderate  excess  of  acid,  and  stir.  If  this  fails  to  dis- 
solve the  carbonate,  add  more,  and  so  on.  The  excess  of 
acid  is  then  neutralized  by  the  alkali,  and  this  excess  may  be 
calculated  from  the  known  ratio  of  the  two  solutions.  This  is 
an  indirect  process  of  standardization. 

7.  Anhydrous  sodium  carbonate  may  be  substituted  for 
calcium  carbonate  in  the  standardization  of  the  acid  solution, 
if  the  purity  of  the  carbonate  can  be  assured.  It  has  the  ad- 
vantage of  solubility,  and  is  better  for  use  with  sulphuric  acid 
since  the  calcium  carbonate  may  become  partially  coated  with 
insoluble  calcium  sulphate,  and  fail  to  be  acted  upon  by  the 
acid. 

Instead  of  standardizing  the  acid  solution  as  described,  it 
is  equally  practicable  to  standardize  the  alkali  solution  against 
purified  oxalic  acid  (C2H2O4.2H2O),  potassium  acid  oxalate 
(KHC2O4.H2O),  potassium  tetroxalate  (KHC2O4.C2H2O4. 
2H2O),  or  potassium  acid  tartrate  (KHC4H4O6).  The  oxalic 
acid  and  the  oxalates  should  be  specially  prepared  to  insure 
purity,  the  main  difficulty  lying  in  the  preservation  of  the 
water  of  crystallization. 

It  should  be  noted  that  the  acid  oxalate  and  the  acid  tar- 
trate each  contain  one  hydrogen  atom  replaceable  by  a  base, 
while  the  tetroxalate  contains  three  such  atoms,  and  the  oxalic 


8 o  VOL UME TRIG  ANAL  YSIS. 

acid  two.     Each  of  the  two  salts  first  named  behave  as  mono- 
basic acids,  and  the  tetroxalate  as  a  tribasic  acid. 

8.  The  sulphuric  acid  solution  may  also  be  standardized 
by  precipitating  the  acid  as  barium  sulphate,  which  is  weighed 
as  in  gravimetric  analysis.  Hydrochloric  acid  may  be  simi- 
larly standardized  by  precipitation  and  weighing,  as  silver 
chloride. 


DETERMINATION     OF     THE     TOTAL    ALKALINE    STRENGTH     OF 
SODA    ASH. 

Procedure.  —  Weigh  out,  on  the  laboratory  balances,  two 
portions  of  soda  ash  of  about  5  grams  each.  Place  them  on 
small  watch-glasses  and  dry  them  at  no0  C.,  until  the  weight 
is  constant  within  0.005  grams,  weighing  at  intervals  of 
thirty  minutes  and  cooling  in  a  desiccator.  Finally  weigh 
accurately,  and  transfer  the  soda  ash  to  a  No.  4  lipped 
beaker ;  weigh  the  watch-glasses,  and  take  the  difference  as 
the  weight  of  the  sample  for  analysis.  Dissolve  the  ash  in 
75  cc.  of  water,  warming  gently,  and  filter  off  the  insoluble 
residue ;  wash  the  filter  until  the  washings  are  freed  from 
carbonate,  cool  the  filtrate  J;o  the  temperature  of  the  labora- 
tory, and  transfer  it  to  a  250  cc.  measuring  flask,  washing 
out  the  beaker  thoroughly.  Add  distilled  water  until  the 
lowest  point  of  the  meniscus  is  level  with  the  graduation  on 
the  neck  of  the  flask,  and  remove  any  drops  of  water  that 
may  be  on  the  neck  above  the  graduation,  with  a  strip  of 
filter  paper ;  make  the  solution  thoroughly  uniform,  by  pour- 
ing it  out  into  a  dry  beaker  and  back  into  the  flask  several 
times. 

Measure  off  50  cc.  of  this  solution  in  a  measuring  flask, 
but  first  pour  into  the  smaller  flask  at  least  two  small  por- 
tions of  the  soda-ash  solution,  and  shake  to  displace  any 
water.  Finally  fill  the  flask  to  the  mark  with  the  solution, 
and  remove  any  drops  on  the  neck  above  the  graduation. 
Empty  it  into  a  beaker  and  wash  out  the  small  flask,  unless 
it  is  graduated  to  deliver  50  cc.  (Compare  note  3,  p.  72.) 
Dilute  the  solution  in  the  beaker  to  100  cc.,  add  two  drops 
of  methyl  orange  solution,  and  titrate  for  the  alkali  with  the 


SODA  ASH.  8 1 

half-normal  acid,  using  the  half-normal  alkali  to  complete  the 
titration,  as  described  under  the  preparation  of  solutions. 
From  the  corrected  volumes  of  acid  and  alkali  employed, 
and  the  data  derived  from  the  standardization,  calculate  the 
percentage  of  alkali  present,  assuming  it  all  to  be  present  as 
sodium  carbonate. 

It  is  advisable  to  measure  out  a  second  portion  of  50  cc. 
from  the  main  solution  in  each  case,  to  confirm  the  original 
titration. 

Notes.  —  i.  Soda  ash  is  crude  sodium  carbonate.  If  made 
by  the  ammonia  process,  it  may  contain  also  sodium  chloride 
sulphate,  and  hydrate ;  when  made  by  the  Le  Blanc  process, 
it  may  contain  sodium  sulphide,  silicate,  and  aluminate,  and 
other  impurities.  Some  of  these,  notably  the  hydrate,  com- 
bine with  acids  and  contribute  to  the  total  alkaline  strength, 
but  it  is  customary  to  calculate  this  strength  in  terms  of 
sodium  carbonate ;  /'.  ^.,  as  though  no  other  alkali  were 
present. 

2.  In  order  to  secure  a  sample  which  shall  represent  the 
average  value  of  the  ash,  it  is  well  to  take  at  least  5  grams. 
As  this  is  too  large  a  quantity  for  convenient  titration,  an  ali- 
quot portion  of  the  solution  is  measured  off,  representing  one 
fifth  of  the  entire  quantity. 

It  is  also  possible  to  weigh  out  exactly  2.6525  grams  of  the 
soda  ash,  dissolve,  filter,  and  titrate  the  entire  solution,  when 
the  number  of  cubic  centimeters  of  half-normal  acid  used  will 
indicate  directly  the  alkaline  strength  in  terms  of  sodium  car- 
bonate. The  student  should  verify  this  statement  by  calcu- 
lation. 

3.  It  is  customary  to  dry  the  soda  ash  at  110°  C.  before 
analysis.     Greater  uniformity  of  results  is  then  attainable. 

Complete  expulsion  of  the  moisture  would  require  a  tem- 
perature just  below  the  fusion  point  of  the  carbonate. 

4.  The    residue   insoluble   in   water  must   be  completely 
washed  to  remove  soluble  alkali,   and  the  filtrate  must   be 
cooled  to  a  temperature  approximating  17.5°  C.  before  dilu- 
tion to  a  definite  volume,  in  a  measuring  flask  which  is  cor- 
rect at  that  temperature. 

A  50  cc.  pipette  may  be  equally  well  employed  to  measure 
out  the  aliquot  portion,  in  place  of  the  50  cc.  flask. 


8 2  VOL UME  TRIG  ANAL  YSIS. 

5.  The  determination  of  the  caustic  alkali  in  the  soda  ash 
may  be  accomplished  by  precipitating.,  the  carbonate  with 
barium  chloride,  removing  it  by  rapid  filtration,  and  titrating 
for  the  alkali  in  the  filtrate.  The  carbonated  alkali  is  then 
calculated  as  the  difference  between  the  caustic  alkali  and  the 
total  alkali. 


DETERMINATION    OF    THE   ACID    STRENGTH    OF    OXALIC   ACID. 

Procedure.  — Weigh  out  two  portions  of  the  acid  of  about 
I  gram  each.  Dissolve  these  in  50  cc.  of  warm  water,  filter, 
if  the  solution  is  not  clear,  and  wash  the  filter  completely 
with  hot  water,  until  freed  from  oxalic  acid.  Add  to  the  fil- 
trate two  drops  of  phenolphthalein  solution,  and  run  in  alkali 
from  the  burette  until  the  solution  is  pink ;  add  acid  from 
the  other  burette  until  the  pink  is  just  destroyed,  and  then 
add  0.5  cc.  (not  more]  in  excess.  Heat  the  solution  to  boil- 
ing for  three  minutes.  If  the  pink  returns  during  the  boil- 
ing, discharge  it  with  acid  and  again  add  0.5  cc.  in  excess 
and  repeat  the  boiling.  If  the  color  does  not  then  reappear, 
add  alkali  until  it  does,  and  a  drop  or  two  of  acid  in  excess 
and  boil  again  for  one  minute.  If  no  color  reappears  during 
this  time,  complete  the  titration  in  the  hot  solution,  ending 
with  the  change  from  colorless  to  pink.  From  the  corrected 
volume  of  alkali  required  to  react  with  the  oxalic  acid,  calcu- 
late the  weight  of  the  latter  present,  in  terms  of  the  crystal- 
lized acid  (C2H2O4.2H2O),  and  from  this,  the  percentage 
purity  of  the  sample. 

Notes. —  i.  It  has  already  been  pointed  out  that  it  is  desir- 
able to  employ  the  same  indicator  throughout  standardization 
and  analysis,  a  statement  which  is  applicable  in  this  instance. 
The  student  is  advised,  if  practicable  within  the  time  devoted 
to  the  subject,  to  re-standardize  acid  and  alkali  using  phe- 
nolphthalein as  the  indicator,  and  concluding  each  titration 
by  boiling  the  solution  as  described  above,  or,  what  is  per- 
haps better,  to  standardize  the  alkali  solution  against  one  of 
the  substances  named  in  note  7,  page  79,  using  phenol- 
phthalein. 


OXALIC    ACID.  83 

The  differences  resulting  from  the  change  of  indicator  are 
small,  and  the  student  may  neglect  them  in  this  instance,  per- 
forming the  titration  as  outlined  above  to  gain  some  practice 
with  phenolphthalein  ;  but  it  should  be  remembered  that  if 
the  highest  accuracy  is  desired,  a  re-standardization  through- 
out is  essential. 

2.  All  commercial  caustic  soda  contains  some  carbonate,, 
and  as  phenolphthalein  is  acted  upon  by  carbon  dioxide,  the 
solution  must  be  boiled  to  expel  the   gas.     Phenolphthalein 
does  not  show  an   alkaline    reaction   with  acid   carbonates ; 
hence  solutions  containing  carbonates  and  this  indicator  be- 
come colorless  when  half  the  carbonate  has  been  acted  upon 
by  acid.     Upon  boiling,  the  bicarbonate  loses  carbon  dioxide, 
forming  normal  carbonate,  and  the  pink  returns.     This  must 
be  again  discharged,  and  the  solution  boiled,  and  so  on. 

A  similar  procedure  is  necessary  with  all  the  indicators 
mentioned,  except  methyl  orange ;  the  latter  does  not,  how- 
ever, give  reliable  results  with  organic  acids,  as  its  own  acid 
is  stronger  than  many  of  them. 

It  is  possible  to  remove  the  carbonate  from  the  caustic 
alkali  by  the  cautious  addition  of  barium  hydroxide  be- 
fore standardization  ;  the  barium  carbonate  is  removed  by 
filtration. 

3.  Hydrochloric  acid  is  volatilized  from  aqueous  solutions, 
except  such  as  are  very  dilute.     If  the  directions  in  the  pro- 
cedure are  strictly  followed,  no  loss  of  acid  need  be  feared, 
but  the  amount  added  in  excess  must  not  exceed  0.5  cc. 

4.  The  end-point  should  be  the  faintest  visible  shade  of 
color,  as  the  same  difficulty  would  exist  here  as  with  methyl 
orange,  if  an  attempt  were  made  to  match  shades  of  pink. 


84  VOLUMETRIC   ANALYSIS. 

II.    OXIDATION   PROCESSES. 

GENERAL    DISCUSSION. 

Under  this  section  may  properly  be  included,  beside  ox- 
idation processes,  the  rather  limited  number  of  methods 
in  which  standard  solutions  of  reducing  agents  only  are 
employed,  since  a  corresponding  oxidation  of  the  reducing 
agent  must  always  be  a  part  of  the  reaction.  In  general  it 
may  be  stated  that  oxidizable  substances  are  determined  by 
direct  titration,  while  oxidizing  substances  are  usually  deter- 
mined by  indirect  methods. 

Many  quantitative  determinations  are  made  possible  by 
the  application  of  these  volumetric  methods  involving  oxida- 
tion or  reduction,  which  are  not  practicable  by  gravimetric 
procedures.  A  notable  example  is  that  of  the  determination 
of  iron  in  the  presence  of  aluminum.  The  mixture  of 
hydroxides  thrown  down  by  ammonia  may  be  ignited, 
weighed,  and  the  iron  subsequently  determined  volumetric- 
ally,  by  solution  of  the  precipitate,  reduction  of  the  iron, 
and  oxidation  by  potassium  permanganate  or  bichromate. 
The  aluminum  is  then  determined  by  difference. 

The  important  oxidizing  agents  employed  for  volumetric 
solutions  are  potassium  bichromate,  potassium  permanga- 
nate, potassium  ferricyanide,  iodine,  ferric  chloride,  and 
sodium  hypochlorite. 

The  important  reducing  agents  which  are  used  in  the 
form  of  standard  solutions  are  ferrous  sulphate  (or  ferrous 
ammonium  sulphate),  oxalic  acid,  sodium  thiosulphate,  stan- 
nous  chloride,  and  arsenious  acid.  Other  reducing  agents? 
as  sulphurous  acid,  sulphuretted  hydrogen,  and  zinc  (nascent 
hydrogen)  take  part  in  the  processes,  but  not  as  standard 
solutions. 

The  most  important  combinations  among  the  foregoing 
are  the  following :  Potassium  bichromate  and  ferrous  salts ; 
potassium  permanganate  and  ferrous  salts  ;  potassium  per- 
manganate and  oxalic  acid  ;  iodine  and  sodium  thiosulphate ; 
hypochlorites  and  arsenious  acid. 


BICHROMA  TE  PROCESS.  85 

BICHROMATE    PROCESS    FOR  THE   DETERMINATION 

OF  IRON. 

GENERAL     DISCUSSION. 

Ferrous  salts  may  be  promptly  and  completely  oxidized  to 
ferric  salts,  even  in  cold  solution,  by  the  addition  of  potas- 
sium bichromate,  provided  sufficient  free  acid  is  present  to 
combine  with,  and  hold  in  solution,  the  ferric  iron  resulting 
from  the  oxidation. 

The  free  acid  may  be  either  hydrochloric  or  sulphuric, 
but  the  former  is  usually  preferred,  since  it  is  by  far  the 
best  solvent  for  iron  and  its  compounds.  The  reaction  in 
the  presence  of  hydrochloric  acid  is  as  follows ;  6  FeCl2 
+  K9Cr907  +  14  HC1  =  6  FeCl3  +  2  CrC)3  +  2  KC1 
+  7  H26. 

The  weight  of  potassium  bichromate  necessary  for  a  nor- 
mal solution  was  shown  on  page  65  to  be  49.08  grams.  It  is 
possible  to  prepare  a  standard  solution  of  the  bichromate  by 
directly  weighing  the  requisite  quantity,,  and  dissolving  it  in 
a  definite  quantity  of  water.  The  commercial  salt,  though 
rarely  sufficiently  pure  for  this  purpose  as  it  often  contains 
potassium  sulphate,  may  be  purified  by  re-crystallization 
from  hot  water,  but  must  then  be  dried,  and  finally  heated 
to  fusion  to  expel  the  last  traces  of  moisture.  The  fusion 
temperature  must  not  be  exceeded,  as  the  bichromate  will 
lose  oxygen  if  more  strongly  heated. 

It  is,  perhaps,  a  more  satisfactory  mode  of  procedure  to 
make  up  a  solution  of  the  commercial  salt  and  determine  its 
strength  by  comparison  with  iron,  either  in  the  form  of  iron 
wire  of  known  purity,  or  ferrous  ammonium  sulphate, 
FeSO4.  (NH4)2SO4.  6  H2O.  A  standard  wire  is  now  offered 
in  the  market  which  answers  the  purpose  well,  and  its  iron 
contents  may  be  determined  for  each  lot  by  a  number  of 
gravimetric  determinations.  It  may  be  best  preserved  in 
jars  containing  calcium  chloride  but  this  must  not  be  al- 
lowed to  come  in  contact  with  the  wire.  It  should,  how- 
ever, even  then  be  carefully  examined  for  rust,  before  use. 


."86  VOLUMETRIC   ANALYSIS. 

If  ferrous  ammonium  sulphate  is  used  as  the  standard, 
clear  crystals  only  must  be  selected,  and  it  is  perhaps  even 
better  to  determine  by  gravimetric  methods,  once  for  all,  the 
iron  contents  of  a  sample  which  has  been  ground  and  mixed. 

It  is  well  to  have  on  hand  a  solution  of  this  salt  for  use  in 
connection  with  the  bichromate.  Such  a  solution  is  more 
stable  if  about  5  cc.  of  concentrated  sulphuric  acid  per 
liter,  are  added. 

It  is  plain  that  all  of  the  iron  in  the  solution  must  be  in 
the  ferrous  condition  before  titration.  The  available  agents 
for  the  reduction  of  any  ferric  iron  are  stannous  chloride, 
sulphurous  acid,  sulphuretted  hydrogen,  and  zinc ;  of  these 
stannous  chloride  acts  most  readily,  the  completion  of  the 
reaction  is  most  easily  noted,  and  the  excess  of  the  reagent 
is  most  readily  removed.  The  latter  object  is  accomplished 
by  oxidation  by  means  of  mercuric  chloride  added  in  excess, 
as  the  mercury  salts  have  no  effect  upon  ferrous  iron  or 
the  bichromate.  The  reactions  involved  are  2  FeCl3 
+  SnCl2  =  2,FeC]2  +  SnCl4;  and  SnCl2  +  2  HgCl2  = 
SnCl4  -f-  2  HgCl.  The  mercurous  chloride  is  precipitated. 

It  is  essential  that  the  solution  should  be  cold  and  that 
the  stannous  chloride  should  not  be  present  in  great  excess, 
lest  a  secondary  reaction  take  place,  resulting  in  the  reduc- 
tion of  the  mercurous  chloride  to  metallic  mercury  :  SnCl2 
-f-  2  HgCl  =  SnCl4  +  2  Hg.  The  occurrence  of  this 
secondary  reaction  is  indicated  by  the  darkening  of  the  pre- 
cipitate, and  since  potassium  bichromate  oxidizes  this  mer- 
cury slowly,  solutions  in  which  it  has  been  precipitated  are 
worthless  as  iron  determinations. 

No  indicator  has  been  found  which  may  be  used  within 
the  solution,  on  account  of  the  deep  green  of  chromium 
salts  ;  the  use  of  potassium  ferricyanide  outside  the  solu- 
tion is  therefore  necessary  to  detect  the  presence  of  fer- 
rous iron.  A  drop  of  the  iron  solution  and  one  of  the 
indicator  solution  are  brought  together  on  a  white  surface 
—  best  a  porcelain  tile  —  and  examined  for  a  blue  precipi- 
tate of  the  ferrous  ferricyanide.  It  is  plain  that  the  potas- 
sium ferricyanide  must  contain  no  ferrocyanide,  for  which 
it  should  be  carefully  tested.  The  latter,  if  present,  may 


BICHROMATE  PROCESS.  87 

be  oxidized  by  the  addition  of  a  little  bromine  to  the  solu- 
tion, after  which  the  ferricyanide  must  be  recrystallized. 

The  indicator  solution  must  be  very  dilute  to  diminish 
the  interference  of  its  own  color  ;  a  crystal  the  size  of  the 
head  of  a  pin  in  25  cc.  of  water  is  an  ample  quantity. 
This  solution  must  be  freshly  prepared  each  day,  as  the 
dilute  solution  is  not  stable. 

No  metals  may  be  present  in  the  titrated  solution  which, 
like  iron,  are  reduced  by  stannous  chloride  and  oxidized  by 
potassium  bichromate  :  notably,  copper,  antimony,  and  plat- 
inum. For  this  reason  a  platinum  crucible  must  never  be 
allowed  to  remain  in  an  iron  solution,  as  ferric  chloride 
exerts  a  slight  solvent  action  upon  the  platinum. 

The  bichromate  solution  may  be  placed  in  burettes  with 
rubber  tips  without  danger  of  deterioration  of  the  solution. 

STANDARDIZATION     OF    A    POTASSIUM    BICHROMATE    SOLUTION. 

Procedure.  —  Pulverize  about  5  grams  of  potassium  bi- 
chromate, dissolve  it  in  water,  and  dilute  to  approximately 

IOOO   CC. 

Pulverize  about  40  grams  of  ferrous  ammonium  sulphate, 
dissolve  it  in  water,  dilute  to  about  1000  cc.,  and  add  5 
cc.  of  concentrated  sulphuric  acid.  It  is  not  necessary  to 
select  clear  crystals  of  the  sulphate. 

Shake  the  solutions  until  they  are  uniform,  and  place 
them  in  burettes,  with  the  precautions  mentioned  on  page 

77- 

Prepare  a  solution  of  potassium  ferricyanide  of  the 
strength  recommended  above,  and  place  single  drops  of 
this  solution  on  the  surface  of  a  porcelain  tile. 

Run  out  from  a  burette  into  a  beaker  about  40  cc.  of  the 
ferrous  solution,  add  15  cc.  of  hydrochloric  acid  (sp.  gr. 
1.12),  dilute  it  to  150  cc.,  and  run  in  the  bichromate  solution 
from  another  burette.  Since  both  the  solutions  are  approx- 
imately deci-normal,  35  cc.  of  the  bichromate  solution  may 
be  added  without  testing.  Test  at  that  point,  by  removing 
a  very  small  drop  of  the  iron  solution  on  the  end  of  a  stir- 
ring rod,  and  mixing  it  with  a  drop  of  indicator  on  the  tile. 


88  VOLUMETRIC   ANALYSIS. 

If  a  blue  precipitate  appears  at  once,  0.5  cc.  of  the  bichnx 
mate  solution  may  be  run.  out  before  testing  again.  The 
stirring  rod  which  has  touched  the  indicator  should  be  dipped 
in  distilled  water,  before  returning  it  to  the  iron  solution. 
As  soon  as  the  blue  appears  to  be  less  intense,  add  the 
bichromate  solution  in  small  portions,  finally  a  single  drop 
at  a  time,  until  the  point  is  reached  at  which  no  blue  color 
appears  after  the  lapse  of  thirty  seconds  from  the  time  of 
mixing  solution  and  indicator  on  the  tile,  the  time  being 
accurately  noted.  At  the  close  of  the  titration  a  large  drop 
of  the  iron  solution  should  be  taken  for  the  test.  To  deter- 
mine the  end-point  beyond  any  question,  as  soon  as  the  thirty 
seconds  have  elapsed  remove  another  drop  of  the  solution 
of  the  same  size  as  that  last  taken,  and  mix  it  with  the 
indicator,  placing  it  beside  the  last  previous  test.  If  this 
last  previous  test  shows  a  blue  tint  in  comparison  with  the 
fresh  mixture,  the  end-point  has  not  been  reached ;  if  no 
difference  can  be  noted  the  reaction  is  complete.  Should 
the  end-point  be  overstepped,  more  ferrous  ammonium  sul- 
phate solution  may  be  added. 

From  the  corrected  volumes  of  the  solutions  used,  cal- 
culate the  value  of  the  ferrous  solution  in  terms  of  the  ox- 
idizing solution. 

Weigh  out  two  portions  of  iron  wire,  of  about  0.24—0.26 
grams  each,  examining  the  wire  carefully  for  rust.  It 
should  be  handled  and  wiped  with  wash-leather,  not  touched 
by  the  fingers,  should  be  weighed  on  a  watch-glass,  and  be 
so  bent  as  not  to  interfere  with  the  movement  of  the 
balance. 

Dissolve  the  wire  in  a  covered  beaker  in  30  cc.  of  hydro- 
chloric acid  (sp.  gr.  1.12),  warming  gently;  wash  off  the 
cover  and  sides  of  the  beaker,  and  add  stannous  chloride 
solution  to  the  hot  liquid,  from  a  dropper,  until  the  solution 
is  colorless,  but  avoid  more  than  a  drop  or  two  in  excess. 
Dilute  with  150  cc.  of  water,  and  cool  completely.  When 
cold,  add  rapidly  about  30  cc.  of  mercuric  chloride  solu- 
tion (50  grams  per  liter).  Allow  the  solutions  to  stand 
three  minutes  and  then  titrate  without  delay,  as  just 
described.  Calculate  the  volume  of  the  bichromate  solution 


BICHROMATE    PROCESS.  89 

which  would  be  required  if  the  solution  were  deci-normal, 
and  add  about  this  quantity.  The  ferrous  ammonium  sul- 
phate solution  may  be  used  if.  the  end-point  is  passed,  and 
much  time  saved. 

From  the  corrected  volumes  of  the  bichromate  solution 
required  to  oxidize  the  iron  actually  present  in  the  wire,  cal- 
culate the  value  of  each  cubic  centimeter  in  terms  of  iron 
(Fe),  and  also  the  amount  of  iron  in  each  cubic  centimeter 
of  the  ferrous  ammonium  sulphate  solution.  Record  these 
values  in  the  notebook. 

Repeat  the  standardization  until  the  results  for  the  value 
of  each  cubic  centimeter  in  terms  of  iron  are  concordant 
within  o.ooooi  gram. 

Notes.  —  i.  Note  carefully  the  statements  in  the  "General 
Discussion  "  bearing  upon  the  steps  in  the  standardization. 

2.  The  hydrochloric  acid  is  added  to  the  ferrous  solution 
to  insure  the  presence  of  sufficient  free  acid  for  the  titration. 

3.  The  time  interval  for  the  indicator  tests  must  be  care- 
fully noted.     Some  time  must  elapse  before  the  completion  of 
the  reaction  when  the  ferrous  iron  is  nearly  all  oxidized,  but  if 
left  too  long,  the  combined  effect  of  light  and  dust  causes  a 
reduction  and  a  deposition  of  a  blue   precipitate,   as  in  the 
process  of  "  blue-printing."     Thirty  seconds  is  a  necessary 
and  also  a  sufficient  interval. 

4.  The   accuracy  of  the  work  may  be  much  impaired  by 
the  removal  of  unnecessarily  large  quantities  of  solution  for 
the  tests.     At  the  beginning  of  the  titration,  while  much  fer- 
rous iron  is  still  present,  the  end  of  the  stirring  rod  need  only 
be  moist  with  the  solution,  but  at  the  close  of  the  titration 
drops  of  considerable  size  may  properly  be  taken  for  the  final 
tests.     The  stirring  rod  should  be  washed  to  prevent  transfer 
of  indicator  to  the  main  solution. 

If  the  end-point  is  determined  as  prescribed,  it  can  be  as 
accurately  fixed  as  that  of  other  methods,  and  if  a  ferrous 
solution  is  at  hand,  the  titration  need  consume  hardly  more 
time  than  that  of  the  permanganate  process,  to  be  described 
later  on. 

5.  The  solution  should  be  allowed  to  stand  about  three 
minutes  after  the  addition  of  mercuric  chloride  to  permit  the 


90  VOLUMETRIC   ANALYSIS. 

complete  deposition  of  mercurous  chloride.  It  should  then 
be  titrated  without  delay  to  avoid  re-oxidation  of  the  iron  by 
the  oxygen  of  the  air. 

6.  The  potassium  bichromate  solution  may  be  diluted  to 
form  an  exactly  deci-normal  solution  if  desired,  according  to 
the  principles  stated  on  page  78. 


DETERMINATION    OF    IRON    IN    LIMONITE. 

Procedure.  —  Grind  the  mineral  to  a  fine  powder,  weigh 
out  two  portions  of  about  0.5  gram  into  No.  7  porcelain 
crucibles,  heat  these  crucibles  to  dull  redness  for  ten  min- 
utes, allow  them  to  cool,  and  place  them,  with  their  contents, 
in  beakers  containing  30  cc.  of  hydrochloric  acid  (sp.  gr. 
1.12).  Heat  at  a  temperature  just  below  boiling  until  the 
undissolved  residue  is  white,  or  until  solvent  action  has 
ceased  ;  if  a  dark  residue  remains,  collect  it  on  a  filter, 
wash  free  from  hydrochloric  acid,  and  ignite  the  filter  in 
a  platinum  crucible.  Mix  the  ash  with  a  small  quantity  of 
sodium  carbonate  and  heat  to  fusion ;  cool,  and  dissolve 
the  fused  mass  in  boiling  water  in  the  crucible.  Unite 
solution  and  precipitate  (if  any)  with  the  acid  extraction, 
washing  out  the  crucible ;  heat  the  solution  to  boiling,  add 
stannous  chloride  solution  until  it  is  colorless,  avoiding  a 
large  excess,  cool,  and  when  cold,  add  40  cc.  of  mercuric 
chloride  solution,  and  proceed  with  the  titration  as  already 
described  on  page  87. 

Calculate  the  percentage  of  iron  (Fe)  in  the  limonite. 

Notes. —  i.  Limonite  is  a  native,  hydrated  oxide  of  iron. 
It  frequently  occurs  in  or  near  peat  beds,  and  contains  more 
or  less  organic  matter,  which,  if  brought  into  solution,  would 
be  acted  upon  by  the  potassium  bichromate.  This  organic 
matter  is  destroyed  by  roasting.  Since  a  high  temperature 
tends  to  lessen  the  solubility  of  ferric  oxide,  the  heat  should 
not  be  raised  above  low  redness. 

2.  A  white  residue,  or  one  known  to  be  free  from  iron, 
may  be  neglected,  and  need  not  be  filtered  off. 


BICHROMATE    PROCESS.  91 

3.  The  quantity  of  stannous  chloride  required  for  the  re- 
duction of  the  iron  in  the  limonite  will  be  much  larger  than 
that  added  to  the  solution  of  iron  wire,  in  which  the  iron  was 
mainly  ferrous.    It  should,  however,  be  added  from  a  dropper. 

4.  The  platinum  crucible,  if  used  for  the  carbonate  fusion, 
must  not  be  put  into  the  iron  solution.     A  platinum  crucible 
may  also  be  used  for  the  roasting,  if  preferred,  with  the  same 
precaution  with  respect  to  the  iron  solution. 

5.  It  is  sometimes  advantageous  to  dissolve  a  large  por- 
tion—  say,  5  grams  —  and  to  take  one  tenth  of  it  for  titra- 
tion.     The  sample  will  then  probably  represent  more  closely 
the  average  value  of  the  ore. 


92  VOLUMETRIC   ANALYSIS. 

DETERMINATION   OF   CHROMIUM   IN   CHROME 
IRON   ORE. 

Procedure.  —  Fuse  10  grams  of  borax  to  a  glass  in  a  large 
platinum  crucible.  Grind  the  chrome  iron  ore  to  the  finest 
powder,  and  weigh  into  the  crucible  about  0.5  gram.  Heat 
over  a  Bunsen  burner,  and  stir  with  a  stout  platinum  wire 
until  a  clear  glass  is  formed  ;  transfer  the  crucible  to  the 
blast  lamp,  add  2  grams  of  sodium  carbonate,  and  heat  to 
quiet  fusion,  using  the  highest  temperature  attainable. 

Mix  1.5  gram  of  sodium  carbonate  and  1.5  grams  of  potas- 
sium nitrate,  and  divide  the  mixture  into  six  portions.  Pro- 
ject these  portions  into  the  crucible  at  suitable  intervals, 
watching  carefully  after  each  addition  to  prevent  loss  of 
material,  particularly  with  the  last  portions.  Stir  the 
liquid  mass  occasionally,  and  continue  the  heating  until  no 
more  gas  is  evolved. 

Cool  the  crucible,  remove  the  fused  mass  as  described  on 
page  32,  and  place  both  in  a  casserole ;  boil  with  water  until 
the  fused  mass  is  disintegrated,  remove  the  crucible,  wash  it 
carefully,  and  add  sulphuric  acid  to  the  solution  in  slight 
excess.  Run  in  ferrous  ammonium  sulphate  solution  from 
a  burette  until  the  ferrous  iron  is  in  excess,  and  titrate  for 
that  excess  with  the  potassium  bichromate  solution. 

From  the  corrected  volumes  of  the  two  standard  solutions, 
calculate  the  amount  of  iron  oxidized  by  the  chromium  from 
the  chrome  iron  ore,  and  subsequently  the  weight  of  chro- 
mium in  solution  and  the  percentage  in  the  ore. 

Notes. —  i.  Chrome  iron  ore  is  essentially  a  ferrous  chro- 
mite,  or  combination  of  FeO  and  Cr2O3.  The  borax  glass 
dissolves  these  oxides,  and  the  nitrate  oxidizes  them  to  ferric 
oxide  (Fe2O3)  and  chromic  anhydride  (CrO3),  the  latter  com- 
bining with  the  alkali  to  form  sodium  chromate.  The  excess 
of  nitrate  is  decomposed,  leaving  the  alkali  as  oxide,  which 
combines  with  some  of  the  non-volatile  acids  present.  Af- 
ter fusion,  the  chromium  is  in  the  same  state  of  oxidation  as 
in  the  bichromate  solution,  and  reacts  in  the  same  way  with 
the  ferrous  salt. 


CHROME    IRON   ORE.  93 

Each  atom  of  chromium  in  the  ore  is  capable  of  oxidizing 
three  atoms  of  iron :  2  Cr2O3  +  3  O2  =  4  CrO3 ;  6  FeSO4 
+  Na2Cr2O7  +  7  H2SO4  =  3  Fe2(SO4)3  -f  Na2SO4  + 
Cr2(S04)3  +  7  H20. 

2.  After  the  oxidation  of  the  chromium  is  approximately 
complete,  the  effervescence  increases  with  each  addition   of 
the  oxidizing  mixture,  and  greater  care  is  required  to  avoid 
loss.     The   determination  requires   con^.'ant    attention   while 
the  crucible  is  over  the  blast  lamp. 

The  use  of  a  graphite  jacket  to  surround  the  crucible 
makes  easier  the  maintenance  of  a  high  temperature,  but  is 
not  indispensable.  Occasional  stirring  prevents  the  accu- 
mulation of  undecomposed  material  at  the  bottom  of  the 
crucible. 

3.  If  a  standard  solution  of  a  ferrous  salt  is  not  at  hand, 
a  weight  of  iron  wire,  somewhat  in  excess  of  that  required 
if  the  substance  were   pure   chromic  oxide  (Cr2O3),  may  be 
weighed  out,  dissolved  in  acid,  reduced  with  stannous  chlo- 
ride, and  treated  as  in  the  standardization  of  the  bichromate 
solution.     The  excess  of  iron  is  then  determined  by  means 
of  the  latter. 


94  VOLUMETRIC   ANALYSIS. 

PERMANGANATE    PROCESS    FOR    THE    DETERMINA- 
TION OF  IRON. 

GENERAL    DISCUSSION. 

Potassium  permanganate  oxidizes  ferrous  salts,  in  cold, 
acid  solution,  promptly  and  completely  to  the  ferric  con- 
dition, while  in  hot,  acid  solution  it  also  enters  into  a  defi- 
nite reaction  with  oxalic  acid,  by  which  the  latter  is  oxi- 
dized to  carbon  dioxide  and  water. 

The  reactions  involved  are  these : 

10  FeSO4  +  2  KMnO4  +  8  H2SO4  =  5  Fe2(SO4)3  -f  K2SO4 
+  2   MnSO4  +  8  H2O. 

5  C2H2O4(2  H2O)  +  2  KMnO4  +  3  H2SO4  =  K2SO4+2MnSO4 
-f  10  CO2  +  8  H2O  -f  (10  H2O). 

These  are  the  fundamental  reactions  upon  which  its  exten- 
sive use  depends,  but  beside  iron  and  oxalic  acid,  the  per- 
manganate enters  into  reaction  with  tin,  copper,  mercury, 
and  manganese  (the  latter  only  in  neutral  solution),  by  which 
these  metals  are  changed  from  the  lower  to  the  higher  state 
of  oxidation,  and  it  also  reacts  with  sulphurous  acid,  sul- 
phuretted hydrogen,  nitrous  acid,  ferrocyanides,  and  most 
soluble  organic  bodies.  It  should  be  noted  that  it  is  only 
with  oxalic  acid,  among  organic  compounds,  that  there  is 
a  definite  reaction  suitable  for  quantitative  purposes. 

From  the  definition  of  a  normal  oxidizing  solution  (page 
65)  the  normal  solution  of  potassium  permanganate  for  use 
in  the  presence  of  acid  must  contain  31.62  grams  of  the 
reagent.  This  is  seen  from  the  following  considerations : 
Two  molecules  of  the  permanganate  yield,  as  shown  in  the 
equations  above,  one  molecule  of  potassium  salt  correspond- 
ing to  the  oxide  K2O,  and  two  molecules  of  manganous  salt 
corresponding  to  the  oxide  MnO,  leaving  five  available  oxy- 
gen atoms  (2  KMnO4=  K2O.-2  MnO.  O5).  The  five  oxygen 
atoms  furnished  by  the  two  gram-molecules  of  the  perman- 
ganate (316.2  grams)  are  equivalent  to  10  grams  of  hydro- 
gen. Accordingly,  the  normal  solution  must  contain  one 


PERMANGANATE  PROCESS.  95 

tenth  of  two  gram-molecules,  or  one  fifth  of  one  gram-mole- 
cule;  i.  e.,  31.62  grams. 

In  neutral  solution  the  permanganate  is  decomposed  as 
indicated  by  the  equation  2  KMnO4  =  K2O.  2  MnO2.  O3. 
The  normal  solution  for  such  purposes  should  contain  one 
sixth  of  two  gram-molecules  ;  i.  e.,  52.7  grams. 

Potassium  permanganate  is  acted  upon  by  hydrochloric 
acid ;  the  action  is  rapid  in  hot  or  concentrated  solutions 
but  slow  in  cold,  dilute  solutions.  The  use  of  the  perman- 
ganate in  the  presence  of  hydrochloric  acid,  or  its  salts,  may 
therefore  be  attended  by  the  possibility  of  error,  and  it  is 
usually  preferable  to  replace  the  hydrochloric  acid  of  iron 
solutions  by  sulphuric  acid,  before  titration.  This  may  be 
done  by  evaporation  with  an  excess  of  the  latter  until  the 
heavy,  white  fumes  of  sulphuric  anhydride  appear. 

The  greater  solubility  of  iron  compounds  in  hydrochloric 
acid  makes  it  desirable  to  titrate,  if  possible,  directly  in  such 
solutions,  and  experiments  made  with  this  end  in  view  have 
shown  that  in  cold,  dilute  hydrochloric  acid  solutions,  to 
which  considerable  quantities  of  manganous  sulphate  or 
chloride,  and  an  excess  of  phosphoric  acid  have  been  added, 
it  is  possible,  with  practice,  to  obtain  satisfactory  results ; 
but  the  end-point  is  less  permanent  than  in  sulphuric  acid 
solutions.  Such  a  process  is  described  in  the  J.  Am.  Chem. 
Soc.y  17,  405.  The  reaction  between  hydrochloric  acid  and 
the  permanganate  is  : 

2  KMnO4  +  16  HC1  =  2  KC1  +  2  MnCl2  +  5C12  +  8  H2O. 

Potassium  permanganate  has  an  intense  coloring  power, 
and,  since  the  solution  resulting  from  the  oxidation  of  the 
iron  and  the  reduction  of  the  permanganate  is  colorless,  the 
latter  becomes  its  own  indicator.  The  slightest  excess  is 
indicated  with  great  accuracy  by  the  color  of  the  solution, 
which  renders  the  titration  one  of  the  most  satisfactory 
known. 

The  commercial  salt  is  rarely  sufficiently  pure  to  admit  of 
direct  weighing  to  form  a  standard  solution,  but  it  may  be 
purified  by  re-crystallization.  The  more  common  practice  is 


96  VOLUMETRIC   ANALYSIS. 

to  standardize  the  solution,  which  may  be  accomplished  by 
comparison  with  iron  wire,  ferrous  ammonium  sulphate,  ox- 
alic acid,  potassium  tetroxalate,  or  potassium  acid  oxalate. 
Other  substances  have  been  proposed,  but  the  foregoing  are 
those  in  common  use. 

The  remarks  on  page  85  referring  to  the  use  of  iron  wire 
and  ferrous  ammonium  sulphate  apply  with  equal  force  here. 
The  pure  oxalic  acid,  or  the  oxalates,  must  be  freshly  pre- 
pared, and  with  great  care ;  they  are  likely  to  lose  water  of 
crystallization  on  standing.  It  must  also  be  borne  in  mind 
that  the  reaction  with  the  oxalates  takes  place  only  in  hot 
solution. 

The  reducing  agents  available  for  the  necessary  reduction 
of  the  iron  before  titration,  are  zinc,  sulphurous  acid,  or  sul- 
phuretted hydrogen ;  stannous  chloride  is  excluded  unless 
the  titration  is  to  be  made  in  the  presence  of  hydrochloric 
acid.  Since  the  excess  of  both  the  gaseous  reducing  agents 
can  only  be  expelled  by  boiling,  with  consequent  uncertainty 
regarding  the  re-oxidation  of  the  iron,  zinc  is  the  more  satis- 
factory agent ;  but  for  prompt  and  complete  reduction  it  is 
essential  that  the  solution  should  be  brought  into  intimate 
contact  with  the  zinc.  This  is  brought  about  by  the  use  of 
a  modified  Jones  reductor,  as  shown  in  the  figure,  page  97.* 

To  prevent  needless  consumption  of  the  zinc,  it  is  first 
amalgamated  by  dissolving  5  grams  of  mercury  in  25  cc.  of 
concentrated  nitric  acid,  diluted  with  an  equal  bulk  of  water, 
and  pouring  into  this  solution  (diluted  to  250  cc.  in  a  1000 
cc.  flask),  500  grams  of  granulated  zinc,  20—30  mesh.  The 
whole  is  shaken  thoroughly  for  two  minutes,  the  solution 
poured  off,  and  the  zinc  washed  thoroughly.  It  may  then  be 
preserved  in  bottles. 

The  tube  A  has  an  inside  diameter  of  18  mm.,  and  is  300 
mm.  long ;  the  small  tube  has  an  inside  diameter  6  mm.,  and 
extends  100  mm.  below  the  stopcock.  At  the  base  of  the 
tube  A  is  coiled  a  piece  of  stout  platinum  wire  ;  on  this  is 
placed  a  plug  of  glass  wool  about  8  mm.  thick,  and  upon 


*  The  details  of  the  reductor  and  preparation  of  the  zinc  are  taken  from 
Blair1  s  Chemical  Analysis  of  Iron,  page  95  et  seq. 


PERMANGANATE  PROCESS. 


97 


this  a  thin  layer  of  asbestos,  such  as  is  used  for  Gooch 
filters,  i  mm.  thick.  The  tube  is  then  filled  with  the  amal- 
gamated zinc  to  within  50  mm.  of  the  top,  and  on  the  zinc 
is  placed  a  plug  of  glass  wool.  The  60  mm.  funnel  B  is 
fitted  into  the  tube 
with  a  rubber  stop- 
per, and  the  reduc- 
tor  is  connected  with 
a  suction  bottle,  F. 
The  bottle  D  is  a 
safety  bottle  to  pre- 
vent contamination 
of  the  solution  by 
water  from  the 
pump. 

The  iron  solution 
during  its  passage 
through  the  reduc- 
tor  comes  into  in- 
timate contact  with 
the  zinc,  and  is  re- 
duced by  the  nascent 
hydrogen  evolved. 
The  column  of  zinc 
should  never  be  less 
than  .  five  inches  in 
length.  Great  care 
must  be  used  to  pre- 
vent the  access  of 
air  to  the  reductor  after  it  has  been  washed  out  ready  for 
use.  If  air  enters,  hydrogen  peroxide  forms,  which  reacts 
with  the  permanganate,  and  the  results  are  worthless. 

It  is  also  possible  to  reduce  the  iron  by  treatment  with 
zinc,  in  a  flask  from  which  air  is  excluded,  the  zinc  being 
completely  dissolved.  This  method  is,  however,  less  con- 
venient and  more  tedious  than  the  reductor. 

Potassium  permanganate  solutions  are  not  usually  stable 
for  long  periods,  and  change  more  rapidly  when  first  pre- 
pared than  after  standing  some  days.  This  is  probably 


9 8  VOLUMETRIC  ANALYSIS. 

caused  by  interaction  with  the  organic  matter  contained 
in  all  distilled  water,  except  that  distilled  from  an  alkaline 
permanganate  solution.  The  solutions  should,  however,  be 
protected  from  the  light  as  far  as  possible,  since  sunlight 
induces  decomposition,  with  a  deposition  of  manganese  diox- 
ide, and  it  has  been  recently  shown*  that  the  decomposition 
proceeds  with  considerable  rapidity,  with  the  evolution  of 
oxygen,  after  the  dioxide  has  begun  to  form.  As  commer- 
cial samples  of  the  permanganate  are  likely  to  be  contami- 
nated by  the  dioxide,  it  is  advisable  to  filter  solutions  through 
.asbestos  before  standardization.  Such  solutions  are  rela- 
tively stable. 

The  permanganate  solution  cannot  be  placed  in  burettes 
with  rubber  tips,  as  a  reduction  takes  place  upon  contact 
with  the  rubber.  The  solution  has  so  deep  a  color  that 
the  lower  line  of  the  meniscus  cannot  be  detected ;  readings 
must  therefore  be  made  from  the  upper  edge. 

STANDARDIZATION    OF    A    POTASSIUM    PERMANGANATE 
SOLUTION. 

Procedure.  —  Dissolve  about  3.25  grams  of  potassium  per- 
manganate crystals  in  200  cc.  of  water,  in  a  beaker,  warm- 
ing to  hasten  solution.  Filter  through  a  layer  of  asbestos, 
cool,  dilute  to  about  1000  cc.,  and  mix  thoroughly.  Fill  a 
glass-stoppered  burette  with  this  solution,  observing  usual 
precautions,  and  fill  a  second  burette  with  the  ferrous  am- 
monium sulphate  solution  prepared  for  use  with  the  po- 
tassium bichromate.  Run  out  into  a  beaker  about  40  cc.  of 
the  iron  solution,  add  10  cc.  of  dilute  sulphuric  acid  (i  15), 
and  run  in  the  permanganate  solution  to  a  slight  perma- 
nent pink.  Repeat,  until  the  ratio  of  the  two  solutions  is 
fixed. 

Weigh  out  into  beakers  two  portions  of  iron  wire  of  about 
0.25  gram  each.  Dissolve  these  in  dilute  sulphuric  acid  (5 
cc.  of  concentrated  acid  and  100  cc.  of  water),  warming  to 
promote  solution.  Meanwhile  prepare  the  reductor  for  use 


*  Morse,  Hopkins,  and  Walker,  Am.  Chem.J.,  18,  401. 


.PERMANGANATE    PROCESS.  99 

as  follows :  Connect  the  suction  bottle  with  the  vacuum 
pump,  fill  the  reductor  while  the  stopcock  is  closed  (or 
nearly  so)  with  warm  dilute  sulphuric  acid  (5  cc.  of  acid  in 
100  cc.  water),  and  then  open  the  stopcock  so  that  the  acid 
runs  through  slowly.  Continue  to  pour  in  acid  until  200 
cc.  have  passed  through,  then  close  the  cock  while  a  small 
quantity  of  liquid  is  still  left  in  the  funnel.  Remove  the  fil- 
trate, and  again  pass  through  100  cc.  of  the  warm,  dilute 
acid.  Test  this  with  the  permanganate  solution.  A  single 
drop  should  color  it  permanently ;  if  it  does  not,  repeat  the 
washing.  Be  sure  that  no  air  enters  the  reductor. 

Pour  the  iron  solution  while  hot  (but  not  boiling)  through, 
the  reductor  at  a  rate  not  exceeding  50  cc.  per  minute. 
Wash  out  the  beaker  with  dilute  sulphuric  acid,  and  fol- 
low the  iron  solution,  without  interruption,  with  175  cc.  of 
the  warm  acid,  and  finally  with  75  cc.  of  distilled  water,, 
leaving  the  funnel  partially  filled.  Remove  the.  filter  bottle, 
and  cool  the  solution  under  the  water  tap.  Add  10  cc.  of 
dilute  sul-phuric  acid,  and  titrate  to  a  faint  pink  with  the 
permanganate  solution  directly  in  the  filter  bottle.  Should 
the  end-point  be  overstepped,  the  ferrous  ammonium  sul- 
phate solution  may  be  added.  From  the  volume  of  the 
solution  required  to  oxidize  the  iron  in  the  wire,  calculate 
the  value  of  each  cubic  centimeter  in  terms  of  metallic  iron 
(Fe).  The  results  should  be  concordant  within  o.ooooi  gram. 

Notes.  —  i.  A  careful  study  of  the  "  General  Discussion" 
should  be  made  in  connection  with  the  steps  of  this  and  the 
following  procedure. 

2.  The  funnel  of  the  reductor  must  never  be  allowed  to 
empty,  and  if  it  is  left  partially  filled,  the  reductor  is  ready 
for  subsequent  use  without  the  previous  washing.  A  prelim- 
inary test  is  always  a  safeguard  against  error,  however. 

If  more  than  a  small  drop  of  permanganate  solution  is 
required  to  color  100  cc.  of  the  dilute  acid  after  the  reductor 
is  well  washed,  an  allowance  must  be  made  for  the  iron  in 
the  zinc. 

The  rate  of  filtration  of  the  iron  solution  should  not  ex- 
ceed that  prescribed,  but  the  rate  may  be  increased  some, 
what  when  the  wash- water  is  added.  It  is  well  to  allow  tha 


ioo  VOLUMETRIC  ANALYSIS. 


iron  solution  to  run  nearly,  but  not  entirely,  out  of  the  funnel 
before  the  wash-water  is  added.  If  it  is  necessary  to  inter- 
rupt the  process,  the  complete  emptying  of  the  funnel  can 
always  be  avoided  by  closing  the  stopcock. 

It  must  be  borne  in  mind  that  only  the  nascent  hydrogen 
is  efficient  as  a  reducing  agent.  That  which  is  visible  is 
molecular  hydrogen  and  without  influence  upon  the  ferric  iron. 

3.  The  dilute  sulphuric  acid  for  washing  must  be  warmed 
ready  for  use  before  the  reduction  of  the  iron  begins. 

4.  The   end-point    is   more    permanent   in   cold   than   hot 
solutions,  possibly  because  of  a  slight  action  of  the  perman- 
ganate fipon  the  manganous  sulphate  formed  during  the  titra- 
tion.     If  the  solution  turns  brown,  it  is  an  evidence  of  insuf- 
ficient acid,  and  more  should   be   immediately  added.      The 
results  are  likely  to  be  less  accurate  in  this  case,  however, 
as    a    consequence   of   secondary  reactions    between   the  fer- 
rous IJ^ri  and  the  manganese  dioxide  thrown  down. 

5.  ine  potassium  permanganate  may,  of  course,  be  diluted 
and  brought  to   an  exactly  y^  solution  from   the   data  here 
obtained.     The  percentage  of  iron  in  the  iron  wire,  as  estab- 
lished by  gravimetric  methods,  must  be  taken  into  account 
in  the  calculation. 


DETERMINATION    OF    IRON    IN    LIMONITE. 

Procedure.  —  Weigh  out  two  portions  of  the  powdered 
limonite,  roast,  and  bring  them  into  solution  as  described 
on  page  90,  but  dissolve  finally  in  casseroles.  Add  cau- 
tiously to  the  solution  5  cc.  of  concentrated  sulphuric  acid, 
and  evaporate  on  the  steam  bath  until  the  solution  is  nearly 
colorless.  Cover  the  casseroles  and  heat  over  the  flame  of 
the  lamp  until  the  heavy  white  fumes  of  sulphuric  anhy- 
dride are  freely  evolved.  Cool  the  casseroles,  add  ioo  cc. 
of  water,  and  boil  until  the  ferric  sulphate  is  dissolved; 
pour  the  warm  solution  through  the  reductor,  proceed  as 
described  under  standardization,  and  titrate  with  the  per- 
manganate solution  in  the  filter  flask,  using  the  ferrous 
ammonium  sulphate  solution,  if  need  be.  From  the  volume 
of  permanganate  solution  used,  calculate  the  equivalent 
quantity  of  iron  (Fe),  and  the  percentage  in  the  limonite. 


PERMANGANATE    PROCESS.  IQI 

Notes.  —  i.  The  preliminary  roasting  is  probably  neces- 
sary, even  though  the  sulphuric  acid  subsequently  chars  the 
carbonaceous  matter.  Certain  nitrogenous  bodies  are  not 
rendered  insoluble  in  the  acid,  and  would  be  oxidized  by 
the  permanganate. 

2.  The  hydrochloric  acid,  both  free  and  combined,  is  dis- 
placed by  the  less  volatile  sulphuric  acid  at  its  boiling  point. 
The  ferric  sulphate  separates  at  this  point,  since  there  is  no 
water  to  hold  it  in  solution,  and  care  is  required  to  prevent 
bumping. 

3.  The  ferric  sulphate  usually  has  a  silky  appearance,  and 
is  easily  distinguished  from  the  flocculent  silica  which  remains 
undissolved.     A  small  quantity  of  glass  wool  may  be  placed 
in  the  neck  of  the  funnel  to  prevent  the  passage  of  this  silica 
into  the  reductor. 


102  VOLUMETRIC  ANALYSIS. 

DETERMINATION  OF  THE  OXIDIZING  POWER  OF 
PYROLUSITE. 

Pyrolusite,  when  pure,  consists  of  manganese  dioxide. 
Its  value,  as  an  oxidizing  agent  and  for  the  production  of 
chlorine,  depends  upon  the  percentage  of  MnO2  in  the 
sample.  This  percentage  is  determined  by  an  indirect 
method,  in  which  the  manganese  dioxide  is  reduced  and 
dissolved  by  an  excess  of  ferrous  sulphate  or  oxalic  acid, 
and  the  unused  excess  determined  by  titration  with  per- 
manganate. 

Procedure.  —  Grind  the  mineral  in  an  agate  mortar  until 
no  grit  whatever  can  be  detected  when  the  powder  is 
placed  between  the  teeth.  Dry  the  ground  sample  on 
a  watch-glass  at  110°  C.  for  an  hour,  transfer  it  to  a 
stoppered  weighing  tube,  and  weigh  out  two  portions  of 
about  0.5  gram  into  No.  3  beakers.  Calculate  the  weight 
of  oxalic  acid  (C2H2O4.  2  H2O)  required  to  react  with  the 
weights  of  pyrolusite  taken  for  analysis,  assuming  it  to  be 
pure  manganese  dioxide:  MnO2  -)-  C2H2O4(2  H2O)  +  H2SO4 
=  MnSO4  +  2  CO2  +2H2O  +  (2  H2O).  We"igh  out  about 
0.2  gram  in  excess  of  this  quantity  of  pure  oxalic  acid 
into  the  corresponding  beakers,  weighing  the  acid  accu- 
rately, and  recording  the  weight  in  the  notebook.  Cover 
the  beakers,  and  pour  into  each  25  cc.  of  water  and  50  cc- 
of  dilute  sulphuric  acid  (i  15);  warm  the  liquid  gently,  until 
the  evolution  of  carbon  dioxide  ceases.  If  a  residue  re- 
mains which  is  sufficiently  colored  to  obscure  the  end-reac- 
tion of  the  permanganate,  it  must  be  removed  by  filtration. 

Finally,  heat  the  solution  to  a  temperature  just  below 
boiling,  and,  while  hot,  titrate  for  the  excess  of  the  oxalic 
acid  with  potassium  permanganate  solution.  From  the  cor- 
rected volume  of  the  solution  required,  calculate  the  amount 
of  oxalic  acid  undecomposed  by  the  pyrolusite ;  subtract 
this  from  the  total  quantity  of  acid  used,  and  calculate  the 
weight  of  manganese  dioxide  which  would  react  with  the  bal- 
ance of  the  acid,  and  from  this  the  percentage  in  the  sam- 
ple. Consult  Part  V,  page  134. 


PYROLUSITE.  103 

Notes.—  i.  The  success  of  the  analysis  is  largely  depend- 
ent upon  the  fineness  of  the  powdered  mineral.  If  properly 
ground,  solution  should  be  complete  in  fifteen  minutes  or  less. 

2.  The  ground   pyrolusite   is  somewhat  hygroscopic.      It 
should  be  dried  at  a  low  temperature  (110°),  as  a  higher  heat 
tends  to  expel    water  of  constitution   from   hydrated  oxides 
which  may  also  be  present. 

3.  A  moderate  excess  of  oxalic  acid  above  that  required 
to  react  with  the  pyrolusite  is  necessary  to  promote  solution- 
otherwise   the   residual   quantity  of  oxalic   acid   would  be  so 
small   that  the  last  particles  of  the   mineral   would  scarcely 
dissolve.     It  is  also  desirable  that  a  sufficient  excess  of  the 
acid  should  be   present  to  react  with  a  considerable  volume 
of  the  permanganate  solution  during  the  titration. 

4.  Care  should  be  taken  that  the  sides  of  the  beaker  are 
not  overheated,  as  oxalic  acid  would  be  decomposed  by  heat 
alone,    if  crystallization    should   occur   on   the    sides    of   the 
vessel.      Strong  sulphuric   acid    also   decomposes  the   oxalic 
acid.     The  dilute  acid  should,  therefore,  be  prepared  outside 
the  beaker. 

5.  Ferrous    ammonium    sulphate    or     iron    wire    may    be 
substituted   for   the    oxalic    acid.     The   reaction   is   then   the 
following:  2  FeSO4  +  MnO2  +  2  H2SO4  =  Fe2  (SO4)3  -f- 
MnSO4  -f-    2  H2O.       The  excess  of   ferrous  iron   may  also 
be  determined  by  means  of  potassium  bichromate,  if  desired. 

Great  care  is  required  to  prevent  the  oxidation  of  the  iron 
by  the  air,  if  ferrous  salts  are  employed. 

6.  Other  volumetric  processes  may  be  employed  for  this 
determination,  one  of  which  is  outlined  in  the  following  reac- 
tions :    MnO2  +   4HC1  =  MnCl2    +    C12  +  2  H2O;   C12 
+  2  KI  =  I2  +  2  KC1;   I2  +  2  Na2S203  =    Na2S4O6  + 
2  Nal.     The  chlorine  generated  by  the  pyrolusite  is  passed 
into  a  solution  of  potassium   iodide.      The   liberated   iodine 
is  then  determined  by  titration  with  sodium  thiosulphate,  as 
described  on  page  107.    This  is  a  direct  process,  although  it 
involves  three  steps. 

It  is  also  possible  to  absorb  and  weigh  the  liberated  carbon 
dioxide  evolved  during  the  reaction  with  the  oxalic  acid,  and 
from  this  weight  to  find  the  percentage  of  manganese  dioxide 
in  the  sample.  This  is  a  gravimetric  process. 


104  VOLUMETRIC   ANALYSIS. 

IODIMETRY. 

GENERAL    DISCUSSION. 

The  titration  of  iodine  against  sodium  thiosulphate,  with 
starch  as  an  indicator,  may  perhaps  be  regarded  as  the 
most  accurate  of  volumetric  processes.  It  may  be  used 
both  in  acid  and  in  neutral  solutions  to  measure  free  iodine, 
and  the  latter  may,  in  turn,  serve  as  a  measure  of  any 
substance  capable  of  liberating  iodine  from  potassium 
iodide  under  suitable  conditions  for  titration.  For  exam- 
ple :  the  quantity  of  potassium  bromate  in  a  commercial 
sample  of  that  salt  may  be  determined  through  the  follow- 
ing reactions  :  KBrO3  +  6  KI  +  3  H2SO4  =  3  K2SO4 
+  KBr  +  3  I2  +  3  H20,  and  I2  +  2  Na.S.O,  =  Na2S4O6 
+  2  Nal. 

Another  illustration  is  afforded  by  the  process  outlined 
in  note  6,  page  103. 

Iodine  is  an  oxidizing  agent,  and,  as  such,  must  conform 
to  the  same  conditions  as  other  similar  bodies,  with  re- 
spect to  its  normal  solutions.  From  the  equation  SO2  +  I2 
+  H2O  =  SO3  +  2  HI,  it  is  plain  that  126.85  grams  of 
iodine  suffice  to  liberate  the  oxygen  necessary  to  oxidize 
I  gram  of  hydrogen,  and  that,  accordingly,  that  weight  of 
iodine  is  requisite  for  a  normal  solution.  Deci-  and  centi- 
normal  iodine  solutions  are  commonly  used. 

Iodine  acts  as  an  oxidizing  agent  either  through  the  de- 
composition of  water,  in  the  presence  of  an  oxidizable  body, 
as  illustrated  by  the  reaction  As2O3  +  2  I2  +  2  H2O  = 
As2O5  +  4  HI,  or  by  increasing  the  proportion  of  the  neg- 
ative constituent  of  a  compound  through  the  direct  with- 
drawal of  the  positive  component,  as  typified  by  the 
equations :  2  Na2S2O3  +  I2  =  Na2S4O6  +  2  Nal,  and  H2S 
+  I2  =  2  HI  +  S. 

The  tendency  of  the  iodine  to  combine  with  hydrogen  is 
not  sufficient  to  cause  it  to  decompose  water,  unless  some 
body  be  present  which  will  readily  combine  with  the  oxygen 
thus  set  free. 

A  complete  equipment  for  iodimetric  work  requires  solu- 


IODIMETRY.  105 

tions  of  iodine,  sodium  thiosulphate,  potassium  iodide,  and 
starch. 

Commercial  iodine  requires  re-sublimation  before  it  can 
be  regarded  as  sufficiently  pure  to  be  weighed  for  a  standard 
solution.  It  should  be  sublimed  between  watch-glasses, 
after  the  addition  of  potassium  iodide  to  unite  with  any 
chlorine  present  in  combination  with  the  iodine,  and  should 
be  subsequently  dried  over  sulphuric  acid.  It  may  then  be 
dissolved  in  a  stoppered  flask,  in  a  solution  of  potassium 
iodide  (about  18  grams  of  the  iodide  to  12  of  the  iodine), 
and  diluted  to  a  definite  volume. 

Its  solutions  are  decomposed  by  sunlight,  with  the  forma- 
tion of  hydriodic  acid,  and  a  high  temperature  tends  to  vol- 
atilize the  iodine.  They  are  not  stable  for  long  periods, 
and  require  frequent  standardization,  against  arsenious  acid, 
anhydrous  sodium  thiosulphate,  or  standard  solutions  of  the 
latter. 

Iodine  solutions  act  upon  rubber ;  hence  only  burettes 
with  glass  stopcocks  should  be  used. 

Sodium  thiosulphate  (Na2S2O3.  5  H2O)  is  rarely  wholly 
pure  as  sold  commercially,  but  may  be  purified  by  crystalli- 
zation, if  need  be.  The  carbon  dioxide  absorbed  from  the 
air  by  distilled  water  decomposes  the  salt,  with  the  separa- 
tion of  sulphur,  and  if  standard  solutions  are  to  be  pre- 
pared directly,  boiled  water,  which  has  been  cooled  out  of 
contact  with  the  air,  must  be  used. 

Solutions  of  the  thiosulphate  must  be  protected  from 
light  and  heat,  both  of  which  promote  decomposition. 
They  may  be  standardized  against  pure  iodine,  or  —  with 
the  intervention  of  potassium  iodide — against  potassium 
bromate,  potassium  iodate,  or  potassium  bichromate.  The 
reactions  on  page  104  indicate  the  principle  involved. 

It  should  be  noted  that  chlorine  and  bromine  oxidize  the 
thiosulphate  to  sulphate,  while  the  iodine  leads  only  to  the 
formation  of  sodium  tetrathionate,  Na2S4O6. 

Commercial  potassium  iodide  generally  contains  a  small 
quantity  of  iodate,  which,  in  acid  solution,  liberates  iodine, 
as  indicated  by  a  yellow  coloration.  The  reaction  is:  KIO3 
+  5  KI  +  3  H2S04  =  3  K2S04  +  3  I2  +  3  H2O.  The 


io6  VOLUMETRIC    ANALYSIS. 

iodate  is  not  necessarily  uniformly  distributed  through  the 
iodide,  and,  in  order  that  an  accurate  blank  test  for  iodate 
may  be  made,  which  shall  apply  to  each  analysis,  it  is  neces- 
sary to  bring  a  considerable  quantity  of  the  iodide  into  solu- 
tion, and  to  take  a  measured  volume  of  this  solution  for 
each  analysis.  The  strength  is  adapted  to  the  work  in 
hand. 

The  starch  solution,  for  use  as  an  indicator,  must  be 
freshly  prepared.  A  soluble  starch  is  now  obtainable  which 
serves  well,  and  a  solution  of  0.5  gram  of  this  starch  in 
25  cc.  of  boiling  water  is  sufficient.  It  is  ready  for  use 
when  cold,  and  from  I  cc.  to  2  cc.  suffices. 

If  soluble  starch  is  not  at  hand,  potato  starch  may  be 
used.  Mix  about  I  gram  with  5  cc.  of  cold  water  to  a 
smooth  paste,  pour  1 50  cc.  of  boiling  water  over  it,  warm  for 
a  moment  on  the  hot  plate,  and  put  it  aside  to  settle.  De- 
cant the  supernatant  liquid  through  a  filter  and  use  the  clear 
nitrate.  5  cc.  of  this  solution  are  needed  for  a  titration. 

The  solution  of  potato  starch  is  less  stable  than  the  sol- 
uble starch.  The  solid  particles  of  the  starch,  if  not  re- 
moved, become  so  colored  by  the  iodine  that  they  are  not 
readily  decolorized  by  the  thiosulphate. 

The  iodo-starch  blue  is  discharged  by  caustic  alkalies,  or 
normal  carbonates  of  the  fixed  alkalies,  but  not  by  the 
bicarbonates. 

STANDARDIZATION    OF    IODINE    AND    SODIUM    THIOSULPHATE 

SOLUTIONS. 

Procedure. — Weigh  out,  on  the  laboratory  balances,  13 
grams  of  commercial  iodine.  Place  it  in  a  mortar  with  18 
grams  of  potassium  iodide  and  triturate  with  small  portions 
of  water  until  all  is  dissolved.  Dilute  the  solution  to 
1000  cc.* 

Weigh  out  25  grams  of  sodium  thiosulphate,  dissolve  it 
in  water,  and  dilute  to  1000  cc. 


*  It  will  be  found  more  economical  to  have  a  considerable  quantity  of  the 
solution  prepared  by  a  laboratory  attendant,  and  to  have  all  unused  solutions 
returned  to  the  common  stock. 


IOD1METRY.  107 

Place  these  solutions  in  burettes  (the  iodine  in  a  glass- 
stopped  burette),  observing  the  usual  precautions  to  prevent 
dilution.  Run  out  40  cc.  of  the  thiosulphate  solution  into 
a  beaker,  dilute  with  150  cc.  of  water,  add  I  cc.  to  2  cc. 
of  the  soluble  starch  solution,  and  titrate  with  the  iodine 
to  the  appearance  of  the  blue  of  the  iodo-starch.  Repeat, 
until  the  ratio  of  the  two  solutions  is  established. 

(Method  A.) 

Weigh  out,  into  No.  4  beakers,  two  portions  of  0.175- 
0.200  gram  each,  of  pure  arsenious  acid.  Dissolve  in  10  cc. 
of  sodium  hydroxide  solution,  with  stirring.  Dilute  the 
solutions  to  150  cc.  and  add  hydrochloric  acid  until  the 
solution  contains  a  few  drops  in  excess,  and  finally  add  a 
concentrated  solution  of  5  grams  of  sodium  bicarbonate 
(HNaCO3).  Cover  the  beakers  to  avoid  loss.  Add  the 
starch  solution,  and  titrate  with  the  iodine  to  the  appear- 
ance of  the  blue  of  the  iodo-starch,  taking  care  not  to  pass 
the  end-point. 

From  the  corrected  volume  of  the  iodine  solution  used  to 
oxidize  the  arsenious  acid,  calculate  the  quantity  of  iodine 
in  each  cubic  centimeter,  and  its  relation  to  the  normal. 
From  the  ratio  between  the  solutions,  calculate  similar  val- 
ues for  the  thiosulphate  solution. 

(Method  B.) 

Weigh  out  into  No.  4  beakers  two  portions  of  about 
0.150-0.175  gram  of  potassium  bromate,  or  potassium 
iodate.  Dissolve  these  in  50  cc.  of  water  and  add  a  suf- 
ficient volume  of  potassium  iodide  solution  to  furnish  3 
grams  of  the  salt.  Add  to  the  mixture  10  cc.  of  dilute 
sulphuric  acid  (i  :  5),  allow  the  solution  to  stand  for  three 
minutes,  and  dilute  to  1500:. ;  run  in  thiosulphate  solution 
from  a  burette  until  the  color  of  the  iodine  is  nearly  de- 
stroyed, then  add  I  cc.  to  2  cc.  of  starch  solution,  titrate 
to  the  disappearance  of  the  iodo-starch  blue,  and  finally 
add  iodine  solution  until  the  color  is  just  restored.  Make 
a  blank  test  for  the  amount  of  thiosulphate  solution  re- 


io8  VOLUMETRIC  ANALYSIS. 

quired  to  react  with  the  iodine  liberated  by  the  iodate  in 
the  potassium  iodide  solution,  and  deduct  this  from  the 
total  volume  used  in  the  titration. .  From  the  data  obtained, 
calculate  the  weight  of  thiosulphate  in  each  cubic  centimeter 
of  the  solution,  and  its  relation  to  a  normal  solution,  and, 
subsequently,  similar  values  for  the  iodine  solution. 

Notes.  —  i.  The  two  methods  of  standardization  seem  to 
yield  equally  satisfactory  results,  and  the  student  is  advised 
to  try  both.  The  arsenious  acid  and  the  potassium  salts  both 
require  careful  examination  to  establish  their  purity.  The 
former  usually  requires  re-sublimation,  and  the  two  latter 
re-crystallization. 

2.  The   color  of  the    iodo-starch  is  somewhat  less  satis- 
factory in  concentrated  solutions  of  alkali  salts,  notably  the 
iodides.     The  dilution  prescribed  obviates  this  difficulty. 

3.  Arsenious  acid  dissolves  more  readily  in  caustic  alkali 
than  in  the  bicarbonates,  but  the  presence  of  caustic  alkali 
during  the  titration  is  not  admissible.     It  is,  therefore,  de- 
stroyed by  the    addition  of  acid,   and  the    solution    is    then 
made  alkaline  with  the  bicarbonate.     Normal  carbonates  of 
the  fixed  alkalies  cannot  be  used. 

The  reaction  during  titration  is  the  following : 

Na8AsO3  +  I2  +  2  HNaCO3  =  Na3AsO4  +  2  Nal 
+  2  C02  +  H20. 

As  the  reaction  between  sodium  thiosulphate  and  iodine 
is  not  always  free  from  secondary  reactions  in  the  presence  of 
even  the  weakly  alkaline  bicarbonate,  it  is  best  to  avoid  the 
addition  of  any  considerable  excess  of  iodine.  Should  the 
end-point  be  passed  by  a  few  drops,  the  thiosulphate  may  be 
used  to  correct  it. 

4.  The  potassium  iodide  should  be  measured  from  a  stock 
solution  for  the  reasons  stated  on  page  105.    It  is  then  pos- 
sible to  make  an  accurate  blank  test  for  the  iodate. 

DETERMINATION    OF    ANTIMONY    IN    STIBNITE. 

The  sample  for  analysis  should  be  pure,  leaving,  at  most, 
only  a  siliceous  residue. 


IODIMETRY.  109 

Procedure.  —  Weigh  out  two  portions  of  about  0.35-0.40 
gram  of  the  mineral  (which  should  be  well  ground),  into  two 
small  dry  beakers  (No.  2).  Pour  over  the  stibnite  5  cc.  of 
hydrochloric  acid  (sp.  gr.  1.20)  and  warm  gently,  but  keep 
well  below  the  boiling  point.  When  the  residue  is  white, 
add  to  each  I  gram  of  solid  tartaric  acid.  Dilute  the  solu- 
tion very  cautiously  by  adding  water  in  portions  of  5  cc., 
stopping  as  soon  as  the  solution  turns  red.  It  is  possible 
that  no  coloration  will  appear,  in  which  case  cautiously  con- 
tinue the  dilution  to  125  cc.  If  a  red  precipitate  or  colora- 
tion does  appear,  warm  the  solution  until  it  is  colorless  and 
again  dilute  cautiously.  Continue  this  to  a  total  volume  of 
125  cc.,  and  boil  for  a  minute. 

Meanwhile,  dissolve  6  grams  of  sodium  bicarbonate  in 
200  cc.  of  water,  in  a  No.  6  beaker ;  pour  the  cold  acid  so- 
lution of  the  antimony  into  this,  avoiding  loss  by  efferves- 
cence. Make  sure  that  the  solution  contains  an  excess  of 
the  bicarbonate,  and  then  add  i  cc.  to  2  cc.  of  starch  solu- 
tion and  titrate  with  iodine  solution  to  the  appearance  of 
the  blue,  avoiding  an  excess. 

From  the  corrected  volume  of  the  iodine  solution  required 
to  oxidize  the  antimony,  calculate  the  weight  of  the  latter  in 
the  solution  and  the  percentage  in  the  stibnite. 

Notes. —  i.  The  success  of  this  determination  is  largely 
dependent  upon  close  adherence  to  the  directions  as  given, 
particularly  with  respect  to  the  amounts  of  reagents  and  the 
dilution. 

2.  Antimony  chloride  is  volatile  with  the  steam  from  its 
concentrated   solutions ;    hence  these   solutions  must  not  be 
boiled  until  they  have  been  diluted. 

3.  The  separation  of  antimony  oxy-chloride  from  solutions 
of  the  chloride,  on  dilution  with  water,  is  prevented  by  the 
addition  of  the  tartaric  acid. 

4.  Stibnite   is  native   antimony  sulphide,   and  upon  solu- 
tion in  hydrochloric  acid  sulphuretted  hydrogen  is  liberated, 
a  part  of  which  is  absorbed  by  the  acid,  unless  the  heating 
is  long   continued.       Upon   dilution,   a   point   is   reached  at 
which  the  sulphide  of  antimony,  being  no  longer  held  in  so- 
lution by  the  acid,  separates.     If  the  dilution  is  immediately 


no  VOLUMETRIC  ANALYSIS. 

stopped  and  the  solution  warmed,  this  sulphide  is  again 
brought  into  solution  and  at  the  same  time,  some  of  the  sul- 
phuretted hydrogen  is  expelled.  This  procedure  must  be 
continued  until  the  sulphuretted  hydrogen  is  all  removed, 
since  it  reacts  with  iodine.  (H2S  -\-  12  =  2  HI  -|-  S.) 

If  no  precipitation  of  the  sulphide  occurs,  it  is  an  indica- 
tion that  it  was  all  expelled  immediately  after  solution. 

5.  If,   for   any   reason,    a   white   precipitate    of  the    oxy- 
chloride  separates  during  dilution  (which  should  not  occur  if 
the  directions  are  followed),  it  is  best  to  discard  the   deter- 
mination and  to  start  anew. 

6.  The  reaction  between  the  iodine  and  the  antimony  is 
parallel  with  that  between  iodine  and  arsenious  acid. 


BLEACHING  POWDER.  Ill 

CHLORIMETRY. 

GENERAL    DISCUSSION. 

Under  chlorimetry  are  included  those  processes  by  which 
not  only  free  chlorine,  but  also  bromine,  hypochlorous  and 
hypobromous  acids  are  estimated.  The  reagent  employed 
is  arsenious  acid,  in  bicarbonate  solution.  In  this  weakly 
alkaline  solution  the  reaction  between  chlorine  and  the 
arsenious  acid  is  parallel  with  that  of  iodine. 

A  solution  of  arsenious  acid '  which  has  been  prepared 
from  the  pure  acid  may  be  used  without  standardization, 
and  is  stable  for  long  periods,  but  the  commercial  acid  re- 
quires re-sublimation  to  remove  arsenic  sulphide,  which 
may  be  present  in  small  quantity.  To  prepare  the  solution, 
dissolve  about  5  grams  of  the  powdered  acid,  accurately 
weighed,  in  10  cc.  of  a  concentrated  sodium  hydroxide  so- 
lution, dilute  the  solution  to  300  cc.,  and  make  it  faintly 
acid  with  hydrochloric  acid.  Add  30  grams  of  sodium 
bicarbonate,  and  dilute  the  solution  to  1000  cc.  in  a  meas- 
uring flask.  If  desired,  the  value  of  this  solution  may 
be  checked  by  titration  against  the  standardized  iodine 
solution. 

The  indicator  required  is  made  by  dipping  strips  of  filter 
paper  in  a  starch  solution,  to  which  I  gram  of  potassium 
iodide  has  been  added.  These  strips  are  allowed  to  drain 
and  spread  upon  a  watch-glass.  When  touched  by  a  drop 
of  the  solution,  the  paper  turns  blue  until  an  excess  of  the 
arsenious  acid  has  been  added.  The  paper  must  be  moist 
when  used. 

DETERMINATION    OF    THE    AVAILABLE    CHLORINE    IN    BLEACH- 
ING  POWDER. 

Procedure.  —  Weigh  out  from  a  stoppered  test-tube  into 
a  porcelain  mortar,  about  3.5  grams  of  bleaching  powder. 
Keep  the  mortar  away  from  the  door  of  the  balance  case, 
to  avoid  injury  to  the  balance.  Triturate  the  powder  in  the 
mortar  with  successive  portions  of  water,  until  it  is  well 
ground  and  transferred  to  a  500  cc.  measuring  flask.  Fill 


H2  VOLUMETRIC   ANALYSIS. 

the  flask  to  the  graduation  and  shake  thoroughly.  Meas- 
ure off  25  cc.  of  this  semi-solution  in  a  measuring  flask, 
or  pipette,  observing  the  precautions  named  on  page  80. 
and  the  further  precaution  that  the  liquid  removed  shall 
contain  its  proportion  of  suspended  matter. 

Empty  the  flask  into  a  beaker  and  wash  it  out.  Run 
in  the  arsenious  acid  solution  from  a  burette,  until  no 
further  reaction  takes  place  with  the  starch-iodide  paper 
when  touched  by  a  drop  of  the  solution  of  bleaching  pow- 
der. From  the  volume  of  solution  required  to  react  with 
the  bleaching  powder,  calculate  the  percentage  of  chlorine 
in  the  latter,  assuming  the  titration  reaction  to  be  that  be- 
tween chlorine  and  arsenious  acid.  Note  that  one  twenti- 
eth of  the  total  weight  of  bleaching  powder  enters  into  the 
reaction. 

Notes.  —  i.  Bleaching  powder  may  be  regarded  as  con- 
taining both  calcium  chloride  and  hypochlorite.  Its  effi- 
ciency, when  treated  with  acids,  depends  upon  the  quantity 
of  the  latter  constituent,  since  the  hydrochlorous  acid  yields 
as  bleaching  agents  both  oxygen  and  chlorine.  It  is  cus- 
tomary, however,  to  express  the  value  of  the  bleaching  agent 
in  terms  of  available  chlorine,  as  though  only  that  were  a 
factor  in  its  efficiency.  The  chlorine  present  as  chloride  is, 
of  course,  not  available  for  bleaching  purposes. 

2.  Bleaching  powder  readily  loses  chlorine  on  exposure 
to  the  air,  as  a  result  of  the  absorption  of  carbon  dioxide. 
The  sample  must  be  carefully  protected,  but  even  then  it  is 
rarely  possible  to  obtain  closely  agreeing  results  from  sepa- 
rate samples.     For  technical  purposes,  it  is  usually  sufficient 
to  examine  one  sample.     The  student  should  check  his  re- 
sults by  titrating  two  portions  from  the  500  cc. 

3.  The  powder  must  be  triturated  until  it  is  fine,  other- 
^  J         wise  the  lumps  will  inclose  calcium  hypochlorite,  which  will 

••'.       ^  fail  to  react  with  the  arsenious  acid.     The  clear  supernatant 

liquid  gives  percentages  which  are  below,  and  the  sediment 
percentages  which  are  above  the  average.     The  liquid  meas- 
ured off  should,  therefore,  carry  with  it  its  proper  proportion 
fl  of  the  sediment. 


SULPHOCYANATE   PROCESS.  113 


III.     PRECIPITATION   METHODS. 


SULPHOCYANATE   PROCESS    FOR   THE    DETER- 
MINATION   OF    SILVER. 


GENERAL    DISCUSSION. 

The  addition  of  a  solution  of  potassium  or  ammonium 
sulphocyanate  to  one  of  silver  in  nitric  acid,  causes  a  depo- 
sition of  silver  sulphocyanate,  as  a  white  curdy  precipitate. 
If  ferric  nitrate  is  also  present,  the  slightest  excess  of  the 
sulphocyanate,  over  that  required  to  combine  with  the  sil- 
ver, is  indicated  by  the  deep  red  which  is  characteristic  of 
the  sulphocyanate  test  for  iron, 

The  reactions  involved  are:  AgNO3  +  KSCN  =  AgSCN 
+  KN03,  and  3  KSCN  +  Fe(NO3)3  =  Fe(SCN)3  + 
3  KN03. 

The  normal  solution  of  the  sulphocyanate  should  contain 
a  sufficient  quantity  of  the  salt  to  combine  with  I  gram  of 
hydrogen  to  form  sulphocyanic  acid  ;  i.  e.,  a  gram-molecule, 
or  97.23  grams  KSCN.  The  sulphocyanate  cannot  be  accu- 
rately weighed  ;  its  solutions  must,  therefore,  be  standard- 
ized against  silver  nitrate,  either  in  the  form  of  a  standard 
solution,  or  by  weighing  out  small  portions.  An  ^  solution 
of  the  sulphocyanate  is  the  strength  to  be  used  for  titrations. 

The  reaction  with  silver  may  be  carried  out  in  nitric  acid 
solution,  and  in  the  presence  of  copper,  if  the  latter  does 
not  exceed  70  per  cent.  Above  that  percentage  it  is  neces- 
sary to  add  silver,  in  known  quantity,  to  the  solution. 

The  liquid  must  be  cold  at  the  time  of  titration  and  en- 
tirely free  from  nitrous  compounds. 

A  saturated  solution  of  ferric  alum,  to  which  a  moderate 
quantity  of  nitric  acid  has  been  added,  serves  as  an  indi- 
cator. The  volume  used  is  5  cc.  and  should  be  the  same  for 
each  titration. 


H4  VOLUMETRIC   ANALYSIS. 

STANDARDIZATION    OF    A    POTASSIUM    SULPHOCYANATE 
SOLUTION. 

Procedure.  —  Crush  a  few  crystals  of  silver  nitrate  in  a 
mortar,  transfer  them  to  a  watch-glass,  and  dry  for  an  hour 
at  110°  C.  Protect  the  nitrate  from  dust  or  organic  matter. 
Weigh  out  two  portions  of  about  0.5  gram  each.  Dissolve 
these  in  50  cc.  of  water  and  add  10  cc.  of  nitric  acid  (sp. 
gr.  1.2),  which  has  been  recently  boiled,  and  5  cc.  of  the 
indicator  solution.  Run  in  the  sulphocyanate  solution  from 
a  burette,  until  a  faint  red  tinge  can  be  detected  in  the  solu- 
tion after  vigorous  stirring.  From  the  corrected  volume 
used,  calculate  the  value  of  the  solution  in  terms  of  metal- 
lic silver,  and  its  relation  to  a  normal  solution.  Repeat, 
until  the  results  are  concordant. 

Notes.  —  i.  The  crystals  of  silver  nitrate  sometimes 
inclose  water,  which  is  expelled  on  drying.  If  the  nitrate  has 
come  into  contact  with  organic  bodies,  it  suffers  a  reduction 
and  blackens  during  the  heating. 

2.  It  is  plain  that  a  standard  solution  of  silver  nitrate 
(made  by  weighing  out  the  crystals)  is  convenient  or  neces- 
sary, if  many  titrations  of  this  nature  are  to  be  made.  In  the 
absence  of  such  a  solution,  the  liability  of  passing  the  end- 
point  is  lessened  by  setting  aside  a  small  fraction  of  the  silver 
solution,  which  can  be  added  at  the  close  of  the  titration  to 
counteract  any  accidental  excess  of  sulphocyanate. 

DETERMINATION    OF    SILVER    IN    COIN. 

Procedure. — Weigh  out  two  portions  of  the  coin,  of  about 
0.5  gram  each.  Dissolve  them  in  15  cc.  of  nitric  acid  (sp. 
gr.  1.2),  and  boil  until  all  the  nitrous  compounds  are  ex- 
pelled ;  cool  the  liquid,  dilute  to  50  cc.,  -add  5  cc.  of  the  in- 
dicator solution,  and  titrate  with  the  sulphocyanate  to  the 
appearance  of  the  faint  red  coloration. 

From  the  corrected  volume  of  the  sulphocyanate  solution 
required,  calculate  the  weight  of  silver  present,  and  the 
percentage  in  the  coin. 

Note.  —  These  solutions,  containing  the  silver  precipitate} 
as  well  as  those  from  the  standardization,  should  be  placed  in 
the  receptacle  for  "  silver  residues." 


PART    IV. 


THE  THEORIES  OF  SOLUTIONS  AND  SOME  OF 

THEIR   APPLICATIONS   TO   ANALYTICAL 

CHEMISTRY.* 


BEFORE  the  development  of  what  may  be  termed  the 
modern  theories  concerning  the  conditions  under  which 
substances  exist  in  solution,  the  exact  nature  of  those 
chemical  reactions  which  are  the  chief  dependence  of  the 
analyst  was  so  little  understood,  that  all  attempts  to  increase 
the  accuracy  of  quantitative  separations  were  made  on  em- 
pirical grounds,  and  were  dictated  solely  by  the  results  of 
experimentation.  With  the  acceptance  and  expansion  of 
these  theories  has  come  the  possibility  of  approaching 
such  problems  from  a  theoretical  standpoint,  and  of  di- 
recting experimentation  along  logical  lines,  the  advantages 
of  which  have  been  most  marked.  In  the  following  pages 
these  modern  conceptions  will  be  briefly  stated;  but  the 
student  is  urged  to  broaden  his  acquaintance  with  the  liter- 
ature bearing  upon  them  as  time  and  interest  permit. 

It  will  be  remembered  that  experiment  has  shown  that  all 
gases,  of  whatever  chemical  character,  show  uniformity  of 
behavior  under  changing  pressure,  obeying  the  law  of  Boyle 
that  the  volume  of  a  specific  quantity  of  a  gas  varies  inversely 
as  its  pressure,  and  that  they  all,  in  obedience  to  the  law 
of  Gay-Lussac,  expand  equally  for  equal  increments  of  tem- 
perature. This  uniformity  of  deportment  is  consistent  only 
with  the  assumption  that  equal  volumes  of  all  gases,  under 
like  conditions  of  temperature  and  pressure,  contain  the 


*In  the  preparation  of  this  chapter  the  author  has  made  frve  use  of  Ostwald1 's 
Scientific  Foundations  of  Analytical  Chemistry,  of  which  he  wishes  to  make  full 
acknowledgment.  The  student  is  referred  to  that  work  for  a  more  extended 
treatment  of  the  topics  than  is  given  in  these  few  pages.  The  brief  outline  here 
given  is  intended  to  serve  chiefly  as  a  foundation  for  later  study,  and  as  an 
incentive  to  further  thought  and  reading.  * 


u6  THEORIES  OF  SOLUTIONS. 

same  number  of  molecules,  or  separate  particles,  and  that 
these  phenomena  are  dependent  only  upon  the  number  of 
these  particles,  and  not  upon  their  chemical  nature.  A 
study  of  gases  bears  out  this  hypothesis,  and  also  the 
assumption  that,  in  general  terms,  these  particles  move 
through  the  space  occupied  by  the  gas  independently  of 
each  other. 

Recent  investigations  have  apparently  furnished  grounds 
for  the  further  hypothesis  that  substances  in  solution  (par- 
ticularly in  dilute  solution)  may  be  compared  with  the  same 
substances  in  the  gaseous  state,  and  that,  if  a  pressure  (the 
existence  of  which  in  a  solution  may  be  experimentally 
proven,  and  which  is  called  osmotic  pressure)  be  substituted 
for  gas  pressure,  the  same  laws  apply  to  the  behavior  of 
these  bodies  in  dilute  solution  as  in  the  gaseous  state.  As 
before,  the  assumption  seems  justified  that  these  physical 
phenomena  may  be  ascribed  to  the  numbers  of  unit  parti- 
cles present,  and  not  to  their  chemical  nature,  and  that  the 
molecules  of  solvent  may  be  regarded  as  without  influence 
upon  the  independence  of  the  movement  of  the  molecules 
•of  the  dissolved  substance,  since  they  act  upon  all  alike. 
But  it  has  also  been  shown  by  experiment  that  when  the 
body  in  solution  is  capable  of  conducting  electricity,  being 
itself  decomposed  at  the  same  time  (that  is,  when  the  sub- 
stance is  an  electrolyte),  there  is  a  larger  number  of  unit 
particles  in  the  solution  than  is  in  accordance  with  the  above 
hypothesis,  and  also  a  larger  number  than  is  found  in  a 
solution  of  an  equivalent  quantity  of  an  indifferent  body, 
or  non-electrolyte.  Such  abnormal  conditions  are  noted 
chiefly  in  aqueous  solutions  of  electrolytes.  .  A  study  of 
the  apparently  abnormal  behavior  of  this  class  of  bodies  led 
Arrhenius  to  propose  his  now  generally  accepted  Theory  of 
Electrolytic  Dissociation,  which  assumes  that  salts,  acids, 
and  bases  —  in  short,  electrolytes  —  in  aqueous  solution  do 
not  exist  solely  in  the  usual  molecular  condition,  but  are 
dissociated  to  a  greater  or  less  degree  into  electrically 
charged  atoms,  or  groups  of  atoms,  called  ions.  The  posi- 
tive ions  (cathions)  are  the  metals,  basic  radicals,  or  hydro- 


THEORIES  OF  SOLUTIONS.  ny 

gen,  on  the  one  hand,  and  the  negative  ions  (anions)  are  the 
acid  radicals,  or  hydroxyl,  on  the  other ;  as,  for  example, 
K  and  Cl  in  potassium  chloride  ;  2  (NH4)  and  SO4  in  ammo- 
nium sulphate ;  H  and  Cl  in  hydrochloric  acid  ;  Na  and  OH 
in  sodium  hydroxide. 

These  ions  are  not  to  be  regarded  as  identical  with  the 
elementary  substances,  sodium,  chlorine,  etc.,  in  the  ordi- 
nary condition  in  which  they  are  more  familiar,  since  each 
ion  bears  a  charge  of  positive  or  negative  electricity,  which 
alters  its  chemical  and  physical  properties  to  a  marked  de- 
gree. In  any  solution  the  amount  of  one  kind  of  electric 
charge  must  be  exactly  equal  to  that  of  the  other  sort ;  for, 
the  moment  that  this  equilibrium  is  disturbed  by  the  dis- 
charge of  any  portion  of  either  sort  of  electricity — as,  for 
example,  when  an  electric  current  is  passed  through  the  so- 
lution,—  the  ions  immediately  pass  to  the  one  or  the  other 
pole,  and  appear  there  with  their  usual  properties.  Thus, 
when  a  solution  of  copper  sulphate  is  submitted  to  the  action 
of  an  electric  current,  the  copper  ions  pass  with  the  positive 
electricity  to  the  negative  pole,  and  there  lose  their  electric 
charge,  and  are  deposited  as  metallic  copper.  In  a  similar 
way  the  SO4  ions  pass  to  the  positive  pole,  where,  when  the 
charge  is  removed,  they  react  with  a  molecule  of  water  to 
form  sulphuric  acid,  with  the  evolution  of  oxygen. 

The  ions  may  be  regarded  as  differing  from  the  corre- 
sponding molecular  bodies  in  the  amount  of  energy  with 
which  they  are  endowed,  and  even  ions  of  identically  the 
same  chemical  composition,  but  of  differing  valency,  may 
exhibit  widely  different  properties.  For  example,  the  ion 
MnO4  in  the  permanganates  (as  KMnO4)  yields  a  purple-red 
solution,  and  behaves  chemically  in  a  way  far  different  from 
the  ion  MnO4  in  the  manganates  (as  K2MnO4),  which  is 
green  in  solution. 

The  degree  of  dissociation  for  specific  dilutions  varies 
with  the  substance ;  it  increases  with  increasing  dilution, 
and  is  theoretically  only  complete  at  infinite  dilution.  The 
degree  of  dissociation  can  be  determined  by  a  study  of 
the  electrical  conductivity  of  the  solutions,  and  by  other 


li 8  THEORIES   OF  SOLUTIONS. 

methods,  which,  however,  do  not  need  to  be  discussed  in 
this  connection.  The  electrolytes  may  be  separated  into 
three  general  classes  with  respect  to  the  degree  of  their 
dissociation:  i.  Those  substances  which  at  moderate  dilu- 
tions are  very  largely  dissociated  (say,  60  to  90  per  cent.) 
—  for  example,  almost  all  salts ;  strong  acids,  like  hydro- 
chloric, sulphuric,  and  nitric  acids ;  and  strong  bases,  as 
the  hydroxides  of  the  alkalies  and  alkaline  earths.  2.  Sub- 
stances of  comparatively  slight  dissociation  (say,  i  to  20  per 
cent.)  at  ordinary  dilution,  as  phosphoric,  sulphurous,  acetic, 
and  most  of  the  organic  acids,  and  ammonia.  3.  Those  sub- 
stances of  which  a  very  small  percentage  only  (less  than 
i  per  cent.)  is  dissociated  at  ordinary  concentrations — as, 
for  example,  weak  acids,  such  as  carbonic,  boracic  or  hydro- 
cyanic acids,  and  hydrogen  sulphide  ;  and  weak  bases,  as  the 
hydroxides  of  di-  and  trivalent  metals.  The  "strength"  of 
an  acid  or  base  corresponds,  then,  in  general  terms,  with  its 
degree  of  dissociation  or  "ionization,"  in  solution. 

The  halogen  compounds  of  cadmium  and  mercury  form 
exceptions  to  the  general  statement  regarding  the  degree  of 
dissociation  of  salts.  These  compounds  are  relatively  little 
dissociated  at  ordinary  dilutions.  No  explanation  of  this 
abnormal  behavior  has  been  offered. 

It  has  also  been  shown  experimentally  that  the  chemical 
properties  exhibited  by  a  solution  of  any  body  are  those  of 
its  ions,  and  that  when  an  element  or  radical  is  not  present 
in  an  ionized  state  those  properties  which  are  considered 
typical,  and  upon  which  reliance  is  placed  for  its  identifica- 
tion or  its  separation,  are  lacking.  For  example,  potassium 
chloride  is  dissociated  into  K  and  Cl  ions  in  aqueous  solu- 
tion, and  the  chlorine  is  easily  detected  by  precipitation  with 
silver  nitrate;  but  neither  chloroform  (CHC13),  nor  potas- 
sium chlorate  (KC1O3)  in  aqueous  solution  show  the  typical 
reaction  for  chlorine,  since  the  former  is  a  non-electrolyte, 
and  the  ions  of  the  latter  are  K  and  C1O3.  The  analytical 
tests  are,  then,  the  tests  for  ions,  and  analytical  chemistry 
is  the  chemistry  of  ions;  and  a  ready  explanation  is  thus 
offered  for  the  fact  that  the  test  for  any  element  or  radical 


THEORIES   OF  SOLUTIONS, 


119 


(for  example,  copper  or  sulphuric  acid)  is  the  same,  no  mat- 
ter what  the  other  element  or  radical  may  be  with  which  it 
was  associated  in  the  molecule,  provided  that  the  molecule 
is  dissociated  in  solution.  This  is  made  apparent  to  the  eye 
by  the  fact  that  the  ions  often  produce  characteristic  colors 
in  solutions.  Thus,  all  cupric  salts  (except  those  in  which 
the  copper  is  associated  with  some  other  colored  ion)  color 
dilute  solutions  blue.  All  chromates  in  dilute  solutions 
are  yellow  and  dichromates  red,  since  the  ion  CrO4  (as  in 
K2CrO4)  is  yellow,  while  the  ion  Cr2O7  is  red  in  solution. 
A  study  of  solutions  with  respect  to  changes  in  color  under 
conditions  of  varying  ionization  has  led  to  useful  results. 
The  applications  in  the  case  of  indicators  used  in  volumetric 
analysis  will  be  made  later  (page  129). 


The  task  of  the  quantitative  analyst  most  frequently  in- 
volves the  separation  from  solution  of  the  substance  to  be 
determined,  in  some  form,  or  in  some  combination  which 
permits  of  quantitative  accuracy,  and  his  chief  aim  is  to 
establish  conditions  which  will  permit  of  such  a  separation. 
Although  it  is  true  that  the  solubility  of  many  bodies  in 
many  solvents  is  less  than  can  be  measured  by  known  means, 
yet,  from  theoretical  considerations,  it  may  be  asserted  that 
no  body  is  absolutely  insoluble  in  any  given  solvent;  and 
it  is  further  true  that  many  substances  employed  for  quan- 
titative separations  are,  even  under  carefully  maintained 
conditions,  appreciably  soluble.  It  is  in  showing  how  this 
solubility  may  be  reduced  to  a  minimum  that  the  develop- 
ment of  and  deductions  from  the  theory  of  Arrhenius  have 
been  of  great  value. 

Take,  for  example,  an  instance  of  quantitative  precipita- 
tion, as  that  of  the  sulphate  ions  from  a  solution  of  ammo- 
nium sulphate  by  means  of  barium  chloride  (page  32).  To 
explain  this  fully  it  is  necessary  to  consider  briefly  the  ques- 
tion of  chemical  equilibrium,  as  it  exists  at  the  moment  when 
a  quantity  of  the  precipitant  has  been  added  which  is  theo- 
retically sufficient  to  complete  the  desired  reaction.  The 


120  THEORIES   OF  SOLUTIONS. 

compounds  in  solution,  being  salts,  are  largely  ionized  into 
the  ions  2  (NH4)  and  SO4,  and  Ba  and  2C1.  Suppose  that 
the  addition  of  the  barium  chloride  were  to  cease  at  the 
instant  when  a  number  of  Ba  ions  has  been  added  exactly 
corresponding  to  the  number  of  (SO4)  ions  present.  The 
compound  BaSO4,  being  relatively  insoluble,  separates  from 
the  solution  for  the  most  part,  but  not  absolutely,  since  a' 
small  quantity  is  still  soluble,  forming  a  saturated  solution  of 
that  sulphate.  The  state  of  equilibrium  which  is  immedi- 
ately produced  is  called  heterogeneous,  because  bodies  in 
different  states  of  aggregation  are  involved.  First,  there  is 
equilibrium  between  the  precipitated  barium  sulphate  and 
that  which  remains  in  solution.  For  such  cases  of  hetero- 
geneous equilibrium  there  must  be,  at  the  contact  surfaces, 
a  definite  ratio  of  concentration  for  the  existing  conditions 
of  temperature.  This  may  be  expressed  by  the  equation 
c'  =  kc,  in  which  c  and  c  are  respectively  the  concentra- 
tions of  the  precipitated  sulphate  and  the  non-ionized  sul- 
phate in  the  solution,  and  k  is  a  constant,  depending  upon 
the  substance  and  the  temperature.  Since,  in  this  instance, 
the  one  body  (the  precipitated  sulphate)  is  a  solid,  its  con- 
centration (c)  does  not  change ;  hence  the  expression  c  =  kc 
simply  states  that,  for  a  given  temperature,  the  concentra- 
tion (c'}  of  the  non-ionized  sulphate  in  solution  is  a  constant 
quantity.* 

Turning  to  the  dissolved  sulphate  and  remembering  that 
as  a  salt  it  is  largely  dissociated,  it  is  found  that  a  second 
condition  of  equilibrium  exists  between  the  undissociated 
and  the  ionized  portions.  The  law  of  mass  action,  the 
truth  of  which  has  been  established  by  experiment,  states 
that  the  chemical  action  of  a  body  is  proportional  to  its 
concentration.  In  the  application  of  this  law  to  electro- 


*The  concentration  of  the  solid  sulphate  is  regarded  as  unchanging,  since  no 
more  of  the  barium  sulphate  remains  in  solution  when  the  precipitate  weighs 
5  grams  than  when  it  weighs  I  gram,  or  even  less.  The  concentration  c'  does 
not  refer  to  the  total  barium  sulphate  in  solution,  but  only  to  the  non-ionized 
portion ;  for,  as  is  seen  from  the  next  paragraphs,  the  undissolved  sulphate  does 
not  stand  in  definite  relations  to  the  concentration  of  the  indvidual  ions,  but  to 
their  product. 


THEORIES  OF  SOLUTIONS.  I2i 

lytes  the  ions  must  be  regarded  as  independent  bodies,  as 
well  as  the  undissociated  molecules  of  the  compound.  The 
application  of  the  law  to  the  case  of  equilibrium  between 
the  ions  Ba  and  SO4,  and  the  non-ionized  BaSO4,  makes  it 
possible  to  express  the  existing  relations  by  the  equation 
ab  =  kfc*  in  which  a  represents  the  concentration  of  the 
Ba  ions,  b  that  of  the  SO4  ions,  and  c  that  of  the  undis- 
sociated BaSO4.  k'  is  again  a  constant  belonging  to  the 
barium  sulphate.  The  product  of  the  concentrations  of  the 
ions  (ab}  is  called  by  Ostwald  the  "solubility-product,"  and 
a  solution  is  saturated  with  a  given  substance  (as  the  barium 
sulphate  in  this  instance)  when  the  product  of  the  concentra- 
tions of  the  ions  is  sufficient  to  yield  this  solubility-product. 
When  it  is  exceeded  the  solution  becomes  supersaturated, 
and  precipitation  should  follow ;  as  long  as  the  value  corre- 
sponding to  saturation  is  not  reached,  more  of  the  substance 
will  pass  into  solution.  This  simple  principle  underlies  the 
entire  question  of  precipitation  from  solutions. 

A  sufficient  number,  then,  of  the  SO4  ions  which  it  is 
desired  to  remove  from  the  solution  have  remained  dissolved 
to  form,  with  the  Ba  ions,  the  solubility-product  of  barium  sul- 
phate. If,  now,  an  additional  quantity  of  barium  chloride  is 
added,  the  concentration  of  the  Ba  ions  is  at  once  increased 
—  i.e.,  the  value  of  a  is  increased;  but  since,  for  constant 
temperature,  c  must  remain  constant,  the  product  ab  must 
also  be  constant,  and  the  concentration  b  must  be  corre- 
spondingly reduced.  This  can  only  happen  as  a  result  of  the 
separation  from  the  solution  of  some  of  the  SO4  ions  in  the 
form  of  barium  sulphate,  and  a  further  addition  of  barium 
chloride  would  cause  further  deposition  of  the  SO4  ions,  and 
so  on,  but  precipitation  could  never  be  absolutely  complete 
as  long  as  any  of  the  solvent  remained ;  it  can,  however,  be 
reduced  to  a  quantity  which  is  negligible. 

The   deduction    to    be    made   from    the   foregoing    is    the 
simple  rule,  that  precipitation  may  be  made  more  complete 


*  Tn  those  instances  in  which  the  substance  is  dissociated  into  ions  of  varying 
valency  the  equation  takes  the  form  anl>m  =  kc,  where  «  and  m  represent  the 
number  of  the  respective  ions. 


I22  THEORIES  OF  SOLUTIONS. 

by  the  addition  to  the  solution  of  some  substance  containing 
an  ion  in  common  with  the  precipitate.  This  is  usually 
effected  by  adding  a  moderate  excess  of  the  precipitant, 
this  excess  to  be  determined  by  considerations  of  expedi- 
ency. In  the  case  cited,  an  undue  excess  of  barium  chloride 
is  to  be  avoided,  on  account  of  the  tendency  of  the  sulphate 
to  drag  out  of  solution  other  bodies  which  may  be  present 
with  it.  (See  page  33,  note  i.) 

It  is  clear  that  the  more  soluble  the  precipitate,  the 
greater  must  be  the  increase  of  the  concentration  of  the 
ion  added,  to  reduce  the  solubility  of  the  precipitated  ion  to 
the  desired  minimum.  The  precipitate  of  magnesium  am- 
monium phosphate  is  a  relatively  soluble  body ;  the  pro- 
portion of  precipitant  required  to  reduce  its  solubility  to 
a  negligible  quantity  is,  therefore,  relatively  large  ;  but,  as 
before,  the  amount  must  be  limited  to  a  reasonable  excess. 

It  should  also  be  noted  that  the  value  of  k  is  large,  and 
of  c  small,  for  precipitates  in  solution  if  they  are  salts,  since 
these  would  be  practically  completely  ionized  at  such  small 
concentrations. 

The  same  principle  is  made  use  of  when  a  wash  water  is 
employed  containing  a  substance  which  has  an  ion  in  com- 
mon with  the  precipitate  treated,  as  in  the  case  of  lead 
sulphate  and  dilute  sulphuric  acid  (page  50).  A  further 
application  of  the  principle  is  made  in  the  preparation  of 
pure  salt,  as  remarked  on  page  22,  note  i,  and  in  the  use  of 
an  ammoniacal  wash  water  for  magnesium  ammonium  phos- 
phate (page  37).  In  the  latter  instance,  the  alcohol  is  added 
simply  as  a  medium  in  which  the  precipitate  is  but  little 
soluble. 

It  must  be  noted  that  some  bodies  tend  to  form  super- 
saturated solutions,  from  which  they  separate  only  slowly 
and  on  long  standing,  as  is  the  case  with  ammonium  phos- 
pho-molybdate.  In  such  solutions  the  solubility-product, 
corresponding  to  equilibrium,  is,  it  is  true,  exceeded,  but  the 
solutions  are  unstable,  and  the  addition  or  formation  of  a 
minute  particle  of  the  dissolved  substance  leads  to  a  sepa- 
ration of  the  solid  until  conditions  consistent  with  saturation 


THEORIES   OF  SOLUTIONS.  123 

are  reached.  The  same  result  is  brought  about  by  vigorous 
shaking  or  stirring,  as  in  the  case  of  the  magnesium  ammo- 
nium phosphate  (page  37). 


It  is  not  always  easy  to  determine  by  inspection  into  what 
ions  a  compound  will  separate  in  solution,  and  the  problem 
is  often  complicated  by  the  intervention  of  secondary  reac- 
tions, by  which  complex  radicals,  which  at  first  act  as  enti- 
ties, subsequently  dissociate  into  simpler  ions.  This  must  be 
borne  in  mind  in  considering  questions  involving  equilibrium. 

It  must  also  be  noted  that  many  precipitates  are  redis- 
solved  to  a  greater  or  less  extent  by  an  excess  of  a  precip- 
itant (in  apparent  contradiction  to  the  principles  just  laid 
down),  as  a  result  of  the  formation  of  double  salts  with 
complex  molecules  or  ions.  This  is  referred  to  again  on 
page  127. 

Again,  it  is  found  that  in  certain  cases  in  which  the 
ions  differ  in  valency,  as  in  dibasic  acids,  the  ionization 

may  take  place  as  follows  :  H2A  =  H  -f-  HA,  and  subse- 
quently HA  =  H  +  A.  As  the  first  hydrogen  ion  behaves 
as  that  of  a  stronger  acid  than  the  second,  and  so  on,  it 
appears  that  these  atoms  differ  somewhat  in  strength. 
This  is  illustrated  by  the  behavior  of  phosphoric  acid 
H3PO4.  The  first  H  ion  is  of  medium  strength,  the  sec- 
ond that  of  a  weak  acid,  while  the  third  can  hardly  be 
replaced. 

Complex  ions  are  formed  from  such  compounds  as  H2PtCl6, 
which  separates  into  2  H  and  PtCl6,  and  K4Fe(CN)6,  which 
yields  4  K  and  Fe(CN)6. 


When  two  or  more  electrolytes  are  brought  together  in 
solution,  the  resulting  phenomena  are  governed  by  principles 
similar  to  those  laid  down  for  the  special  case  of  precipita- 
tion already  cited. 

Two  salts,  such  as  potassium  chloride  and  sodium  nitrate, 
which  furnish  no  ions  capable  of  forming  an  insoluble  body, 


124 


THEORIES  OF  SOLUTIONS. 


exert  no  action  the  one  upon  the  other.  Both  are  largely 
ionized,  and  the  only  bodies  which  could  result  from  an 
interaction,  namely,  sodium  chloride  and  potassium  nitrate, 
are  also  equally  dissociated  into  ions  identical  with  those  of 
the  original  compounds. 

If,  however,  the  ions  are  capable  either  of  uniting  to  form 
a  substance  with  a  smaller  solubility-product  on  the  one 
hand,  or,  on  the  other  hand,  a  compound  which,  under  the 
existing  conditions,  is  less  ionized  than  those  from  which  it 
is  formed,  an  interaction  takes  place  in  that  direction,  and 
the  ions  disappear,  as  such,  to  form  the  non-dissociated  sub- 
stance. If  the  solubility-product  is  sufficiently  small,  a  pre- 
cipitate falls,  as  in  the  case  of  the  barium  sulphate,  while 
the  reaction  between  sodium  hydrate  and  hydrochloric  acid 
may  be  taken  as  an  example  of  lessened  ionization.  The  ions 
in  the  latter  case  are  Na  and  OH,  and  H  and  Cl.  As  water 
is  among  the  least  dissociated  bodies,  the  H  and  OH  ions 
disappear  to  form  non-ionized  water.  (It  is  estimated  that 
it  would  require  1,000,000  liters  of  water  to  furnish  i  gram 
of  hydrogen  ions.) 

The  same  statement  is  true  when  nitric  acid  and  potas- 
sium hydroxide,  or  hydrochloric  acid  and  barium  hydroxide 
react,  from  which  it  is  evident  that  the  net  effect  of  the 
neutralization  of  the  acid  by  the  alkali,  in  each  case,  is  the 
formation  of  nonrionized  water ;  and  the  fact  that  the  heat 
of  neutralization  is  the  same  in  each  of  these  cases,  for 
equivalent  quantities  of  acid  and  alkali,  is  explained. 

Similarly,  if  hydrochloric  acid  is  added  to  a  solution  of 
sodium  acetate  (or  the  reverse),  an  interaction  takes  place, 
whereby  the  H  ions  of  the  hydrochloric  acid  unite  with  the 
C2H3O2  ions  from  the  sodium  salt,  to  form  acetic  acid,  which, 
as  a  weak  acid,  is  relatively  little  dissociated.  The  hydrogen 
ions  of  the  mineral  acid  are  thus  largely  withdrawn  from  the 
solution  in  the  form  of  the  non-ionized  acetic  acid.  But, 
since  acetic  acid  is  far  more  largely  dissociated  than  water, 
this  reaction  will  never  approach  as  near  completion  as  the 
one  previously  mentioned.  With  a  salt  of  a  still  weaker  acid 
the  reaction  would  be  more  complete,  while  the  addition  of 


THEORIES  OF  SOLUTIONS,  125 

hydrochloric  acid  to  a  cold  solution  of  a  nitrate  vthe  two 
acids  being  ionized  to  about  the  same  degree)  produces  no 
effect  with  respect  to  the  ions. 

Similar  principles  may  be  applied  to  the  behavior  of  strong 
and  weak  bases. 

The  "strength"  of  an  acid,  as  has  been  stated,  depends 
upon  the  degree  to  which  it  is  ionized  —  that  is,  upon  the 
concentration  of  its  hydrogen  ions;  and  the  "strength"  of 
a  base  upon  the  concentration  of  its  hydroxyl  ions.  Any 
cause  which  alters  these  concentrations  lessens  or  increases 
the  efficiency  of  the  compound.  If  to  a  solution  of  a  weak 
acid,  a  salt  of  the  same  acid  is  added,  the  strength  of  the 
latter  is  lessened.  For  example,  the  addition  of  sodium 
acetate  to  acetic  acid  makes  of  it  a  weaker  acid.  If  the 
solution  of  the  acid  alone  is  regarded,  the  state  of  equilib- 
rium is  expressed  by  the  equation  ab  =  kc,  where  a  is  the 
concentration  of  the  H  ions,  b  that  of  the  C2H3O2  ions, 
c  that  of  the  undissociated  acid,  and  k  is  a  constant  applying 
to  acetic  acid.  If,  now,  sodium  acetate  is  added  to  the  solu- 
tion, of  which  the  ions  are  Na  and  C2H3O2,  the  concentra- 
tion b  of  the  latter  is  at  once  increased. 

Since  acetic  acid  is  a  substance  which  is  but  little  disso- 
ciated, much  the  larger  part  of  the  acid  is  present  in  the 
non-ionized  condition,  and  the  concentration  (c)  represents 
approximately  that  of  the  entire  quantity  of  acid  present 
(a  and  b  are  relatively  small  in  comparison  with  c).  It  is 
clear,  then,  that  while  c  may  be  slightly  increased,  it  cannot 
grow  to  any  large  extent ;  and  since  b  is  disproportionately 
increased  by  the  addition  of  a  large  number  of  C2H3O2  ions 
from  the  acetate,  a  must  diminish,  and  accordingly  some  of 
the  H  ions  unite  with  the  C2H3O2  ions  to  form  a  small 
additional  quantity  of  undissociated  acid.  That  is,  the 
concentration  of  the  hydrogen  ions,  which  was  small  at 
the  start,  diminishes,  and  the  acid  power  is  weakened.  It 
is  plain  that  the  weaker  the  acid  is  in  the  first  place,  the 
greater  the  relative  diminution  of  strength.  No  perceptible 
change  takes  place  with  the  strong  acids  unless  the  salt  is 
added  in  excessive  quantities.  This  principle  may  fre- 


J26  THEORIES  OF. SOLUTIONS. 

quently  be  applied  in  the  case  of  weak  acids,  when  it  is 
desired  to  lessen  their  solvent  action  in  quantitative  pre- 
cipitations, as  that  of  zinc  sulphide  from  acetic  acid 
solutions. 

Following  the  same  reasoning,  it  is  evident  that  the  addi- 
tion of  a  strong  acid  to  the  solution  of  a  weak  one  lessens 
the  ionization  of  the  latter.  This  is  of  importance  in  a 
study  of  precipitations  by  sulphuretted  hydrogen,  which 
itself  is  a  weak  acid.  An  interesting  example  is  that  of 
the  precipitation  of  zinc  from  its  salts.  If  sulphuretted 
hydrogen  is  passed  into  a  solution  of  zinc  chloride  a 
portion  of  the  zinc  is  thrown  down  as  sulphide,  but  the 
action  soon  ceases.  This  may  be  accounted  for  by  the 
following  considerations :  I.  Zinc  sulphide  can  only  sep- 
arate where  the  product  of  the  concentration  of  the 
Zn  and  S  ions  in  solution  exceeds  its  solubility-product. 

2.  Sulphuretted    hydrogen    is    but    very    slightly    ionized. 

3.  As    the   zinc    sulphide    separates,    hydrochloric    acid    is 
formed  and  is  nearly  completely  ionized.     4.  The  presence 
of  this  strong  acid  (by  increasing  the  concentration  of  the 
H  ions)  gradually  lessens  the   ionization  of  the  weak'  acid, 
sulphuretted  hydrogen,  just  as  sodium  acetate  lessens  that 
of  acetic  acid,  until  the  concentration  of  the  S  ions  becomes 
so  small  that   its  product  with  the  concentration  of  the  Zn 
ions  fails  to  equal  or  exceed  the  solubility-product  of  zinc 
sulphide,  and  precipitation  must  cease.     But  if  zinc  acetate 
is  the  salt  used,  the  whole  of  the  zinc  is  thrown  down  by  sul- 
phuretted hydrogen,  because  the  acetic  acid  liberated  is  not 
a  sufficiently  strong  acid  to  so  far  increase  the  concentration 
of  the  hydrogen  ions,  as  to  cause  a  diminution  of  the  ion- 
ization of  the  sulphuretted  hydrogen  sufficient  to  cause  the 
solubility-product  of  the  zinc  sulphide  to  fall  below  the  crit- 
ical value. 

The  application  of  the  law  of  mass  action  to  the  re-solu- 
tion of  precipitates  may,  after  what  has  been  stated,  be  well 
shown  by  an  example ;  for  instance,  the  solution  of  ferric 
hydroxide  in  hydrochloric  acid.  As  no  body  is  wholly 


THEORIES   OF  SOLUTIONS. 


127 


insoluble  in  water,  a  minute  quantity  of  the  hydroxide  will 
first  pass  into  solution  in  water,  and  in  such  a  solution  this 
substance  will  be  ionized  (slightly,  as  it  is  a  weak  base). 
The  saturated  solution  will  assume,  for  the  moment,  a  con- 
dition of  equilibrium  parallel  to  that  of  barium  sulphate,  as 
stated  on  page  120,  the  ions  being  Fe  and  3  (OH).  But  the 
latter  will  at  once  unite  with  the  H  ions  of  the  hydrochloric 
acid  to  form  water  (undissociated),  and  in  order  to  maintain 
the  concentration  of  the  OH  ions  more  of  the  hydroxide 
must  pass  into  solution.  This  process  repeats  itself  with 
great  rapidity  until  the  hydroxide  is  all  dissolved,  or  the  acid 
saturated.  Similarly,  hydrochloric  acid  will  dissolve  calcium 
oxalate,  because,  by  the  union  of  the  2  H  and  C2O4  ions,  some 
undissociated  oxalic  acid  is  formed,  and  more  oxalate  must 
pass  into  solution  to  replace  the  C2O4  ions  so  removed,  and 
thus  to  restore  the  equilibrium.  It  must  be  noted  here  that 
saturation  of  hydrochloric  acid  with  respect  to  the  oxalate 
will  occur  much  more  quickly  than  with  respect  to  the 
hydroxide,  because  the  oxalic  acid  is  relatively  largely  dis- 
sociated in  comparison  with  water,  and  the  solution  after 
a  time  becomes  saturated  with  Ca  and  C2O4  ions. 

Again,  acetic  acid  dissolves  calcium  phosphate  because 
the  phosphoric  acid  is  less  dissociated  than  acetic  acid ; 
while  calcium  oxalate  is  not  dissolved  because  the  reverse 
is  true.  The  slightly  increased  solubility  of  barium  sulphate 
precipitates  in  the  presence  of  hydrochloric  acid  is  probably 
due,  in  the  same  way,  to  the  lesser  degree  of  dissociation  of 
the  sulphuric  than  the  hydrochloric  acid. 

The  re-solution  of  a  precipitate  is  often  occasioned  by  the 
formation  of  some  complex  ion  in  place  of  the  simple  ions 
formerly  existing.  Thus  copper  salts  are  dissolved  by  so- 
lutions containing  ammonia  with  the  formation  of  complex 
ions,  silver  salts  in  potassium  cyanide,  and  so  on.  This 
must  be  taken  into  account  in  seeking  for  the  explanation 
of  the  solubility  of  precipitates  in  special  cases.  Magnesium 
forms  complex  ions  with  ammonium  compounds,  which  pre- 
vents the  precipitation  of  its  hydroxide  by  ammonia,  a  fact 
which  is  taken  advantage  of  in  qualitative  and  quantitative 
analysis. 


I28  THEORIES  OF  SOLUTIONS. 

It  has  already  been  stated  that  the  water  in  aqueous 
solutions  is  but  very  slightly  ionized.  While  this  is  true, 
it  is  also  a  fact  that  the  H  and  OH  ions  of  the  dissociated 
water  exert  a  distinct  influence  upon  certain  classes  of  salts 
—  namely,  those  in  which  one  or  more  of  the  ions  is  that 
of  a  very  weak  base  or  acid.  For  example,  a  solution  of 
a  neutral  salt,  as  zinc  sulphate,  shows  an  acid  reaction, 
because  the  ions  of  the  salt,  Zn  and  SO4,  together  with 
the  minute  quantity  of  H  and  OH  ions  from  the  dissociated 
water,  tend  to  form  Zn(OH)2  and  H2SO4.  The  former  is 
a  very  weak  base,  and  accordingly  is  but  little  dissociated, 
and  does  not  exhibit  an  alkaline  reaction ;  while  the  sul- 
phuric acid  is,  on  the  contrary,  largely  dissociated,  and  the 
concentration  of  the  H  ions  in  solution  is  sufficient  to  cause 
an  acid  reaction.  The  reverse  process  goes  on  with  a  salt 
of  the  character  of  potassium  cyanide,  from  which  KOH  and 
HCN  are  formed.  The  latter  is  so  weak  an  acid  as  to  be 
but  very  slightly  ionized ;  while  the  potassium  hydroxide 
yields  OH  ions,  which  cause  an  alkaline  reaction.  This 
process,  whereby  the  ions  of  the  water  enter  into  play  with 
the  ions  of  the  substances  in  solution,  is  called  hydrolysis. 

The  precipitation  of  a  basic  carbonate  of  zinc,  mentioned 
on  page  55,  is  the  result  of  hydrolytic  action  of  the  water 
upon  the  zinc  carbonate  first  formed,  by  which  a  portion 
is  changed  to  hydroxide,  which  at  first  is  held  in  colloidal 
solution  (see  page  12),  and  is  then  precipitated  with  the  zinc 
carbonate  upon  boiling. 

It  has  already  been  pointed  out  that  the  action  of  an  acid 
upon  an  hydroxide  is,  in  effect,  merely  the  union  of  H  and 
OH  ions  to  form  water.  Such  a  reaction  remains  incom- 
plete to  an  extent  corresponding  to  the  degree  of  dissocia- 
tion of  the  water ;  but  it  must  be  again  emphasized  that  this 
is  very  small,  and  usually  beyond  the  limits  of  measurement 
permitted  by  the  indicators  in  use.  It  is  a  fact,  however, 
that  the  use  of  certain  indicators  is  disturbed  by  secondary 
hydrolytic  reactions,  as  is  shown  on  page  130. 


THEORIES  OF  SOLUTIONS.  129 

The  indicators  used  in  acidimetry  or  alkalimetry  are  or- 
ganic bodies  either  weakly  basic  or  acid  in  their  character. 
Those  compounds  have  been  selected  as  indicators  which 
furnish  a  colored  ion  when  dissociated,  or  afford  a  colored 
solution  when  in  the  un-ionized  state,  whereby  the  change 
from  the  one  to  the  other  condition  is  made  evident.  For 
example,  methyl  orange  is  a  compound  belonging  to  a  class 
of  organic  bodies  known  as  sulphonic  acids.  In  the  free 
state  the  acid  is  but  very  little  ionized,  and  the  undisso- 
ciated  acid  yields  a  pink  solution.  The  moment  that  the 
acid  is  neutralized  by  alkali,  a  salt  of  the  sulphonic  acid  is 
formed,  which,  in  common  with  all  salts,  is  at  once  disso- 
ciated, and  the  anion  yields  a  yellow  solution.  The  addition 
of  even  a  drop  of  free  mineral  acid  in  excess,  at  once  causes 
a  combination  of  the  H  ions  from  this  acid  and  the  anions 
from  the  salt  of  the  sulphonic  acid,  to  form  the  non-disso- 
ciated sulphonic  acid,  with  the  return  of  the  pink.  This  is 
the  point  chosen  as  the  end-point. 

Phenolphthalein  is  another  weak  acid.  It,  however,  pre- 
sents a  reverse  behavior.  The  undissociated  acid  is  color- 
less in  solution,  while  the  anion  of  the  salt  is  pink.  In  the 
case  of  litmus  the  undissociated  acid  is  red,  the  dissociated 
ion  blue. 

It  is  stated  on  page  72  that  methyl  orange  may  be  used 
in  cold  solutions  in  the  presence  of  carbonic  acid  and  sul- 
phuretted hydrogen,  and  will  still  indicate  the  point  at  which 
all  the  mineral  acid  present  is  neutralized.  The  other  indi- 
cators require  the  previous  expulsion  of  the  carbon  dioxide 
or  sulphuretted  hydrogen,  by  boiling.  The  acid  of  methyl 
orange,  while  still  a  weak  acid,  is  yet  stronger  —  i.  e.,  more 
largely  dissociated  —  than  carbonic  acid,  or  sulphuretted  hy- 
drogen. Accordingly  the  concentration  of  the  H  ions  of 
these  latter  acids  is  less  than  that  of  the  H  ions  from  an 
equivalent  quantity  of  the  sulphonic  acid,  and  the  former  do 
not  enter  into  combination  with  the  OH  ions  of  the  base 
until  after  those  of  the  sulphonic  acid  of  the  indicator  —  /.  e., 
until  after  a  salt  of  the  sulphonic  acid  has  formed  and  the 
end-point  has  been  reached. 


130  THEORIES   OF  SOLUTIONS. 

The  acid  contained  in  phenolphthalein,  litmus,  and  the 
like,  are  weaker  —  /.  e.,  less  dissociated  —  than  carbonic  acid 
or  sulphuretted  hydrogen.  The  latter,  therefore,  first  fur- 
nish H  ions  to  unite  with  the  OH  ions  of  the  hydroxide,  and 
no  salt  of  the  indicator  acid  is  formed  until  these  acids  have 
been  neutralized.  It  is  accordingly  necessary  to  expel  these 
volatile  acids  by  boiling,  in  the  presence  of  an  excess  of 
mineral  acid,  in  order  to  discover  exactly  the  point  at  which 
the  mineral  acid  is  neutralized.  Methyl  orange  is  to  be 
preferred  as  an  indicator  in  alkalimetry,  because  it  forms 
salts  most  readily  with  bases,  and  phenolphthalein  in  acic1- 
imetry,  because  it  responds  most  readily  to  the  influence  of 
acids.  Methyl  orange  is  too  strong  an  acid  to  be  success- 
fully used  with  the  organic  acids  in  general. 

On  the  other  hand,  phenolphthalein  does  not  give  accurate 
results  with  ammonia,  or  in  the  presence  of  ammonium  salts 
(page  73).  Ammonia  is  a  weak  base,  and  the  salt  which 
should  be  formed  with  the  indicator  undergoes  hydrolysis, 
whereby,  instead  of  the  colored  ion,  the  weak,  undissociated 
acid  of  the  indicator  and  undissociated  base  (ammonia)  are 
formed,  and  it  is  only  when  a  relatively  considerable  excess 
of  the  ammonia  exists  in  solution,  that  the  concentration  of 
the  OH  ions  becomes  sufficiently  large  to  overcome  the  hy- 
drolytic  action. 

The  reaction  between  sodium  carbonate  and  a  mineral 
acid  presents  some  interesting  features.  When  sodium 
carbonate  is  placed  in  aqueous  solution  the  carbonic  acid, 
owing  to  its  tendency  to  pass  into  a  less  ionized  state,  gives 
rise  to  the  formation  of  an  ion  HCO3  in  solution;  that  is,  the 
sodium  carbonate  yields  first  the  ions  2  Na  and  CO3,  and  the 
latter,  acting  upon  water,  yield  a  certain  proportion  of  HCO3 
and  OH  ions,  the  latter  occasioning  the  alkaline  reaction. 
When  H  and  Cl  ions  are  added  to  such  a  solution,  in  the 
form  of  hydrochloric  acid,  the  H  and  OH  ions  first  unite  to 
form  water,  and  as  the  OH  ions  disappear,  more  HCO3  ions 
are  formed  from  the  CO3  ions.  This  continues  until  all  the 
CO3  ions  are  transformed  into  HCO3  ions,  when  the  latter 
begin  to  interact  with  the  H  ions  of  the  mineral  acid  to  form 
undissociated  H2CO3. 


THEORIES   OF  SOLUTIONS.  131 

It  is  stated  on  page  81  that  phenolphthalein  may  be  used 
as  an  indicator  to  show  when  one  half  the  carbonate  in  a 
solution  of  an  alkaline  carbonate  has  been  neutralized,  pro- 
vided the  volume  of  water  is  sufficient  to  retain  all  the  lib- 
erated carbonic  acid  in  solution.  From  the  foregoing  it 
appears  that  the  indicator  forms  a  salt  with  the  alkali  as 
long  as  hydroxyl  ions  remain  in  the  solution,  but  as  soon 
as  these  are  withdrawn  to  form  undissociated  water,  the 
hydrogen  ion  of  the  carbonic  acid  unites  with  the  anion  of 
the  phenolphthalein  to  form  the  free,  undissociated  acid. 
As  the  carbonic  acid  is  unstable,  the  carbon  dioxide  tends 
to  escape  even  in  the  cold,  and  is  completely  removed  with 
the  steam  from  boiling  solutions. 

Sodium  bicarbonate  yields  as  ions  Na  and  HCO3,  and 
does  not,  therefore,  give  an  alkaline  reaction  with  phenol- 
phthalein.* Upon  boiling  its  solutions  the  ion  is  decom- 
posed as  follows:  Na  +  HCO3  =  Na  +  OH  +  CO2.  The 
carbon  dioxide  escapes,  and  the  solution  contains  a  quantity 
of  OH  ions  equivalent  to  that  formed  from  one  molecule  of 
normal  sodium  carbonate  for  each  two  molecules  of  the  bi- 
carbonate. This  is  the  explanation  of  the  return  of  the 
color  on  boiling  carbonate  solutions,  when  phenolphthalein 
or  litmus  is  used  as  an  indicator,  as  described  on  page  73. 

As  the  end-points  are  sharper,  the  greater  the  difference 
between  the  degree  of  ionization  of  the  indicator  and  that 
of  the  acid  or  base  used  for  titration,  it  is  better,  in  general, 
to  use  concentrated  rather  than  very  dilute  solutions. 


*It  is  probably  true  that  by  hydrolytic  action  the  NaliCOg  is  slightly 
dissociated  into  Na,  OH,  and  HCO3  ions,  and  that  a  minute  quantity  of  the 
sodium  salt  of  the  indicator  acid  is  formed,  but  so  little  that  the  hydrolytic 
action  dissociates  it.  On  the  other  hand,  the  amount  of  OH  ions  is  sufficient 
to  form  salts  with  other  indicators  (as  litmus)  having  stronger  acids,  the  salts 
being,  therefore,  less  subject  to  hydrolysis. 


PART   V. 


STOICHIOMETRY. 


The  stoichiometrical  problems  with  which  the  analytical  chemist 
has  to  deal  are  not,  as  a  rule,  difficult  either  to  solve  or  to  compre- 
hend. The  student  will  find  that  a  moderate  time  devoted  to  the 
thoughtful  study  of  these  problems  will  do  much  to  prevent  em- 
barrassment in  later  professional  experience,  where  the  ability  to 
make  the  necessary  calculations  for  the  interpretation  of  analyt- 
ical data  is  no  less  important  than  the  manipulative  skill  by  which 
the  data  are  obtained. 

Detailed  solutions  of  a  few  typical  problems  are  given  below. 
The  student  should  study  these  carefully,  and  assure  himself  that 
they  are  fully  understood. 

i.  A  "chemical  factor"  expresses  the  ratio  between  a  spe- 
cific quantity  of  a  chemical  compound  and  the  equivalent  quantity 
of  some  other  body.  For  example,  if  it  is  wished  to  determine 
the  weight  of  sulphur  which  corresponds  to  a  specific  weight  of 
barium  sulphate,  the  latter  is  multiplied  by  the  factor,  or  ratio 

represented  by  the  fraction  BaSQ  ,  or  ^'^  =  °'I373-     It  may 

233-5        32-07 
also  be   expressed  by   the   proportion    BaSO4  :  S  =  wt.  BaSO4 

:  x.  from  which  it  is  plain  that  x  =    3  '°7  .  wt.  BaSO4. 

233-5° 
Again,   if  the   weight  of  FeO   in   Fe2O3   is  desired,   the  factor 

becomes  -^ — =-  =  >      '       =  0.9000.      Similarly,    the   factor   for 
Fe2O3         160.04 

K.2O  94.22 

the   conversion  of    KC1  to  K2O    is    — ^-^   =  -        -   =  0.6320. 

2  K.C1          149.12 

The  logarithmic  equivalents  of  these  values  are  called  log  factors. 
In  the  calculation  of  these  factors,  the  atomic  or  molecular  rela- 
tions of  the  two  substances  must  be  kept  clearly  in  mind  ;  thus,  it 
is  plainly  incorrect  to  express  the  ratio  of  ferrous  to  ferric  oxide  by 

FeO 

the  fraction  ^ — :--,  since  each  molecule  of  the  higher  oxide  must 
re2u3 

correspond  to  two  molecules  of  the  lower.     Carelessness  in  this 
respect  is  one  of  the  most  frequent  sources  of  error. 


S  TO  1C HIO  ME  TRY.  133 

2.  To  calculate  the  volume  of  a  reagent  required  for  a  specific 
operation,  it  is  necessary  to  know  the  exact  reaction  which  is  to  be 
brought  about,  and,  as  with  the  calculation  of  factors,  to  keep  in 
mind  the   molecular  relations  between  the  reagent  and  the  sub- 
stance reacted  upon.      For  example,   to  estimate    the    weight  of 

barium  chloride  necessary  to  precipitate  the  sulphur  from  o.i  gram 

48870 

of  pure  pyrite  (FeS2),  the  proportion  should  stand  2  BaCl2.  2  H2O,: 
120.16 
FeS2  =  x  :  o.i,   where  x  represents  the  weight  of  the  chloride 

required.  Each  of  the  two  atoms  of  sulphur  will  form  a  mole- 
cule of  sulphuric  acid  upon  oxidation,  which,  in  turn,  will  require 
a  molecule  of  the  barium  chloride  for  precipitation.  To  determine 
the  quantity  of  the  barium  chloride  required,  it  is  necessary  to  in- 
clude in  its  molecular  weight  the  water  of  crystallization,  since  this 
is  inseparable  from  the  chloride  when  it  is  weighed.  This  applies 
equally  to  other  similar  instances. 

If  the  strength  of  an  acid  is  expressed  in  percentage  by  weight, 
due  regard  must  be  paid  to  its  specific  gravity.  For  example,  hy- 
drochloric acid  (sp.  gr.  1.12)  contains  23.8  per  cent.  HC1  by 
weight;  /.  <?.,  0.2666  gram. 

3.  No    rules   for    universal    application    to    "  indirect    gravi- 
metric analyses  "  can  be  laid  down.      A  single  example  will  be 
explained. 

Given  a  mixture  of  KC1  -\-  NaCl  weighing  0.15  gram,  which 
contains  53  per  cent,  chlorine,  to  calculate  the  weight  of  KC1  and 
of  NaCl  in  the  mixture. 

The  weight  of  chlorine  in  the  mixture  is  (0.15  X  0-53)  or  0.0795 
gram.  Assuming  that  this  chlorine  was  all  in  combination  with 
potassium,  the  corresponding  weight  of  KC1  would  be  0.1672 
gram  (Cl  :  KC1  =  0.0795  :  0-1672).  This  is  an  excess  of  0.0172 
gram  over  the  actual  weight  of  the  mixture,  and  it  is  plain  that 
this  difference  is  occasioned  by  the  replacement  of  certain  of  the 
molecules  of  potassium  chloride,  weighing  74.56  units,  by  mole- 
cules of  sodium  chloride  weighing  58.50  units.  To  express  this, 
let  it  be  supposed  that  the  mixture  is  made  up  of  n  molecules 

KC1  and  n'  molecules  NaCl ;  then  it  may  be  said  that  n  KC1  -f- 

58.50  74.56  74.56 

n  NaCl  —  0.15  gram,  and  n  KC1  -j-  nr  KC1  =  0.1672  gram,  then 

by  subtracting    the    first   equation  from    the    second  it  is   shown 

74-56  58-5° 

that   n'  (KC1  —  NaCl)  =  0.0172    gram.      That    is,    the    differ- 
ence in  weight  is  equal  to  n'  times  the  difference  in  the  molecu- 


134  $  TO  1C  HI  O  ME  TR  Y. 

lar   weights  of  the   two   chlorides.     The   actual    weight    of   NaCl 

present  (x)  is  equal  to  58.50^,  or,  since  n'  = °'01^2  •,  x  = 

74-S6  —  58-5° 
/        0.0172        \ 
58.50  (         ^  o        ).      This    may   be    expressed    in    the   form 

{74.56  —  58.50)  :  58.50  =  0.0172  :  x,  from  which  x  =.  0.0626. 
The  weight  of  NaCl  subtracted  from  that  of  the  mixture  gives 
the  weight  of  KC1. 

The  weights  of    the  chlorides  may  also  be  calculated  algebra- 

35  4S 
ically  by  solving  the  equations  x  -f-  y  =  0.15    and  — — >  .v  -\- 

jj'      y  =  0.0795,  where  x  is  the   weight  of  KC1   and  y  is   the 

o     "0 

weight  of    NaCl   in   the   mixture. 

4  It  is  sometimes  desirable  to  weigh  out  such  a  quantity  of 
substance  for  analysis,  that  the  number  of  cubic  centimeters  of 
standard  solution  entering  into  the  reaction  shall  represent  directly 
the  percentage  of  the  desired  constituent.  This  may  be  readily 
done,  by  considering  the  relation  of  the  solution  to  a  normal  solu- 
tion and  the  atomic  or  molecular  weight  of  the  desired  component. 
For  example,  suppose  it  is  desired  to  calculate  such  a  weight  for 
K2CO3  in  pearlash,  when  a  half-normal  acid  solution  is  used. 
Since  half-normal  acid  and  alkali  solutions  are  equivalent,  and 
since,  by  definition,  the  half-normal  K2CO3  solution  contains  34.55 
grams  per  liter,  each  cubic  centimeter  of  the  acid  solution  must  be 
equivalent  to  0.03455  gram  K2CO3.  Hence,  100  cc.  would  neu- 
tralize 3.455  grams  pure  K2CO3,  and  this  becomes  the  desired 
weight  of  the  pearlash.  Similarly,  the  required  weight  of  limonite, 
where  the  iron  (Fe)  is  to  be  determined  by  means  of  a  deci-normal 
K2Cr2O7  solution,  is  0.5602  gram. 

5.  One  of  the  most  frequently  recurring  cases  in  volumetric 
analysis  is  that  in  which  it  is  wished  to  express  the  value  of  a  spe. 
cific  solution  in  terms  of  some  substance  other  than  that  against 
which  it  has  been  standardized  ;  as,  for  instance,  the  value  of  a 
permanganate  solution  which  has  been  standardized  against  oxalic 
acid,  in  terms  of  iron.  Although  such  problems  apparently  vary 
widely,  there  are  common  principles  which  can  be  applied  to  them 
all.  These  are  stated  below,  and  the  student  should  assure  him- 
self that  they  are  fully  understood. 

Suppose,  for  example,  it  is  desired  to  find  the  iron  value 
(Fe)  of  a  permanganate  solution,  of  which  i  cc.  is  equivalent  to 
0.0063  gram  C2H2O4.  2  H2O. 


S  TOICHIOME  TRY.  135 

From  a  comparison  of  the  reactions  on  page  94,  it  is  seen  that 
10  molecules  of  ferrous  sulphate  and  5  molecules  of  oxalic  acid 
each  react  with  the  same  amount  (2  molecules)  of  the  perman- 
ganate. These  two  quantities  being,  then,  equivalent  to  the  same 
third  quantity,  must  be  equivalent  to  each  other  ;  in  other  words, 
10  molecules  of  ferrous  sulphate  and  5  molecules  of  oxalic  acid 
have  the  same  reducing  power.  But,  as  stated  above,  the  value  is 
desired  in  terms  of  metallic  iron  (Fe),  not  FeSO4,  but  as  it  is  plain 
that  10  FeSO4  are  equivalent  to  10  Fe,  it  is  proper  to  make  the 

proportion 

560.2  630.25 

10  Fe  :  5  C2H2O4.  2  H2O  =:  x  :  .006302 

in  which  x  —  0.005602  gram.  Here,  again,  as  in  example  2,  it  is 
necessary  to  include  the  water  of  crystallization  in  the  molecular 
weight  of  the  oxalic  acid,  as  it  is  weighed  with  it. 

The  same  conclusion  is  arrived  at,  if  we  consider  the  relation  of 
the  solution  to  the  normal.  As  given,  it  is  deci-normal  and  must, 
therefore,  be  equivalent  to  a  deci-normal  solution  of  iron.  From 
the  equations  cited,  it  is  seen  that  10  FeSO4  unite  with  5  O,  there- 
fore each  molecule  is  equivalent  to  i  hydrogen  atom  in  reducing 
power.  The  normal  solution  must,  then,  contain  i  gram-molecule 
of  ferrous  sulphate,  or  56.02  grams  Fe,  and  each  cubic  centimeter 
of  the  deci-normal  solution  would  contain  0.005602  gram,  the 
value  obtained  above. 

Again,  suppose  the  value  of  the  same  permanganate  solution 
were  desired  in  terms  of  molybdenum  (Mo),  the  reactions  with 
permanganate  being 

5  Mo12Oi9  +   T7  Mn2O7  ;=  60  MoO3  -f-  34  MnO,  and 
5  C2H2O4.  2  H20  +  Mn'2O7  =  2  MnO  +  ioCO2  +  15  H2O. 
(Mn2O7  is  the  anhydride  of  HMnO4.) 

It  is  plain  that  in  these  equations  as  they  stand,  the  molecular 
quantities  of  oxidizing  agent  are  not  equal.  They  can  be  made 
so  by  simply  multiplying  the  second  equation  by  17,  and  they  then 
become, 

5  Mo12Oi9  -f-  17  Mn2O7  =  34  MnO  -|-  60  MoO3,  and 
85  C2H2O4.  2  H2O  +  17  Mn2O7  =  34  MnO  +  170  CO2  + 
255  H20. 

It  is  now  possible  to  reason  in  the  same  way  as  before,  and  to 
conclude  that  85  molecules  of  the  oxalic  acid  have  the  same  reduc- 


136  STOIC 'H '10  ME  TR  Y. 

ing  power  as  5   molecules  of  the  oxide   Moi2Oi9,  or  60  atoms  of 
molybdenum.     Accordingly, 

575.88  10714.25 

60  Mo  :  85  C2H2O4.  2  H2O  =  x  :  0.006302 

in  which  x  =  0.0003387  gram. 

Since  5  Mo12Oi9  unite  with  85  O,  a  normal  solution  of  the 
former  as  a  reducing  agent,  would  contain  T^  of  the  5  gram-mole- 
cules or  3.387  grams  Mo,  and  the  deci-normal  solution  0.3387 
grams  per  liter.  This  agrees  with  the  values  already  obtained. 

6.  It  is  sometimes  necessary  to  calculate  the  value  of  solutions 
according  to  the  principles  just  explained,  when  several  successive 
reactions  are  involved.  Such  problems  may  be  solved  by  a  series 
of  proportions,  but  it  is  usually  possible,  after  stating  these,  to 
eliminate  the  common  factors  and  solve  but  a  single  one. 

For  example,  suppose  it  is  desired  to  express  the  value  of  a 
permanganate  solution,  of  which  i  cc.  =  0.008  gram  iron  (Fe), 
in  terms  of  calcium  oxide  (CaO).  The  reactions  involved  in 
the  volumetric  determination  of  calcium  are  the  following  :  CaCl2 
+  (NH4)2C2O4  =  CaC2O4  +  2  NH4C1  ;  CaC2O4  +  H2SO4 
+  2  H2O  =  CaSO4  +  C2H204.  2  H2O  ;  5  C2H2O4.  2  H2O  + 
2  KMnO4  +  3  H2SO4  =  K2SO4  +  MnSO4  +  10  CO2  + 
18  H2O. 

From  the  considerations  stated  under  5,  the  following  propor- 
tions may  be  made  : 

10  Fe  :  5  C2H2O4.  2  H2O  =  .008  :  x 
5  C2H2O4.  2  H2O  :  5  CaC2O4  =  x  :  y 
5  CaC2O4  :  5  CaO  =  y  :  z 

Canceling  the  common  factors,  there  remains  simply 

560.2  280.4 

10  Fe  :  5  CaO  =  .008  :  z 

Similarly,  from  the  reactions  given  in  note  6,  page  103,  the  equiv- 
alent of  the  iodine  liberated  may  be  calculated  in  terms  of  MnO2 
as  follows  :  Supposing  the  weight  of  iodine  to  be  0.5  gram,  then 

2  I   :  2  KI  =  0.5   :  x 
2  KI   :  2  Cl  —  x  :  y 
2  Cl   :  2  HC1  =  y  :  z 
2  HC1   :  MnO2  =  z  :  w 

Canceling  the  common  factors,   there  remains 
2  I   :  MnO2  =  0.5    :  w 


STOICHIOME  TRY.  137 

To  solve  such  problems  as  5  and  6,  it  is  necessary  to  know  the 
reactions  involved,  and  the  way  in  which  the  various  components 
break  up  ;  then  to  compare  the  reactions  and  to  search  for  those 
molecular  quantities  of  the  compounds  in  question,  which  are 
equivalent  in  their  action  upon  a  common  agent.  Having  found 

these,  as  shown  above,  express  the  molecular  ratio  between  them 

253.7      86.99 
in  the  form  of  a  proportion ;  as,  for  example,  2  I  :  MnO  =  0.5  :  w. 

Expressed  in  the  form  w  =  —     —  0.5,  it  is  plain  that  this  ratio  is 

in  no  way  different  in  principle  from  the  chemical  factor  men- 
tioned in  paragraph  i  ;  indeed,  it  is  the  factor  for  the  conversion 
of  iodine  to  manganese  dioxide. 

PROBLEMS. 

(The  reactions  necessary  for  the  solution  of  these  problems  are 
either  stated  with  the  problem,  or  may  be  found  in  the  earlier 
text.  The  atomic  weights  used  are  those  given  in  the  table  on 
page  145-) 

GRAVIMETRIC    ANALYSIS. 

1.  Calculate    the    chemical    factors    for    (a)    (NH4)2O    from 
2  (NH4)C1.  PtCl4;    (£)    for    K    in    2  KC1.  PtCl4 ;    (V)    for    P    in 
Mg2P2O7 ;    (d)  for  Fe2O3  from  Fe3O4. 

Answers :  (a)  0.1175,  (b)  0.1611,  (c)  0.2787,  (d)  1.034. 

2.  If  0.5  gram  of  platinum  remains  after  the  ignition  of  the 
precipitate   of   the    double   salt,    2  NH4C1.  PtCl4,    derived    from 

1  gram  of  an  ammonium  compound,   calculate  the  percentage  of 
NH3   in   the   latter.      2  NH4C1.  PtCl4    =    2  NH3    +    2  HC1  + 

2  C12  +  Pt.  Answer:  8.76   %. 

3.  What  weight  of  Mn3O4  corresponds  to  i  gram  of  Mn2P2O7  ? 

Answer :  0.5375  gram. 

4.  How  many  cubic  centimeters  of  aqueous  ammonia  (sp.  gr. 
0.96)  containing  9.90  per  cent.  NH3  by  weight,  will  be  required  to 
precipitate   the   iron    as    Fe(OH)3  from    i    gram    of   (NH4)2SO4. 
FeSO4.  6  H2O  ?  Answer:  1.37  cc. 

5.  How  many  cubic  centimeters  of  HNO3  (sp.gr.  1.135)  con- 
taining 20  per  cent.  HNO3  by  weight,  are  required  to  oxidize  the 
iron  in  i  gram  of  FeSO4.  (NH4)2SO4.  6  H2O,  in  the  presence  of 
sulphuric    acid?       6    FeSO4     +     2    HNO3     +     3    H2SO4    — 

3  Fe2  (SO4)3  -f  2  NO  +  4  H2O.  Answer:  0.24  cc. 


138  PROBLEMS. 

6.  The  ignited  precipitate  of  Fe2O3  -f-  A12O3  from   1.5  grams 
of  a  silicate  weighs  0.4069  gram  ;  this  mixture  loses  0.0200  gram 
when   ignited   in    hydrogen.      What    is   the   percentage  of    Fe2O3 
and  A12O3  in  the  sample  ?     Fe2O3  -f  3  H2  =  2  Fe  -(-  3  H2O. 

Answer:  22.68  %  A12O3 ;  4.44  %  Fe2O3. 

7.  How  many  cubic  centimeters  of  "  magnesia  mixture  "  (64 
grams  MgCl2  per  liter)  will  be  required  to  precipitate  the  arsenic 
from  0.2  gram  As2S3,  after  oxidation  to  arsenic  acid?     H3AsO4  -\- 
MgCl2  +  3  NH4OH  =  MgNH4AsO4  +  2  NH4C1  +  3  H2O. 

Answer:  2.42  cc. 

8.  How  many  cubic  centimeters  of  an  ammonium  oxalate  so- 
lution   [(NH4)2C2O4.  H2O]    (40   grams  per  liter),  are  required  to 
precipitate    the    calcium    as    oxalate    from    i    gram     of    apatite 
(Ca3(PO4)2.  CaCl2)  ?     How  many  cubic  centimeters  of  "  magnesia 
mixture  "    (containing    64  grams  MgCl2  per  liter)    are   necessary 
to    combine    with   the    phosphoric    acid   in    the    same    weight    of 
apatite  ?  Answer:  33.73  cc.  and  7.06  cc. 

9.  If  a  calcium  oxalate  precipitate  (which  is  contaminated  by 
silica)  from  0.83  gram  of  dolomite  be  ignited  under  such  condi- 
tions   that    the   decomposition  products  may  be   passed    through 
Ba(OH)2     solution,     and    the     resulting    precipitate    of     barium 
carbonate   be  found,   on   drying,    to  weigh   0.9500  gram,  what   is 
the  percentage  of  CaO  in  the  sample?  Answer:  32.51  %. 

10.  How  many  cubic  centimeters  of  a  potassium  tetroxalate 
solution   (KHC2O4.  C2H2O4.  2  H2O),    containing    50    grams    per 
liter,  would  be  required  to  precipitate  the  calcium  from    i  gram 

'of  a  sample  of  dolomite  yielding  2  %  Fe2O3,  10  %  MgO,  and 
45  °/0  CO2,  assuming  the  iron,  magnesium,  and  calcium  to  be 
present  wholly  as  carbonates,  the  iron  as  ferrous  carbonate  ? 

Answer:   19.05  cc. 

11.  How  many   cubic   centimeters   of    sulphuric   acid   (sp.   gr. 
1.75),  containing  81  per  cent.  H2SO4  by -weight,  are  necessary  to 
replace  the  nitric  acid  in  the  nitrates  formed  from  5  grams  of  a 
brass  containing  65  %  Cu,  34.5  %  Zn  and  0.5  %  Pb  ? 

Answer:  5.37  cc. 

12.  If  5.23  grams  of  brass  yield  0.0345  gram  PbSO4,  and  sub- 
sequently 0.0031  gram  PbO2  on  electrolysis  of  the  filtrate,  what  is 
the  percentage  of  Pb  in  the  brass  ?  Answer :  0.5  %. 

13.  If  in  the  analysis  of  a  brass  containing  65  %   copper,  an 
error  is  made  in  weighing  a  5-gram  portion,  by  which  o.ooi  gram 


S  TOJCH1OME  TR  Y.  139 

too  much  is  weighed  out,  what  would  be  the  percentage  of  copper, 
as  determined  ?  If  the  same  error  is  made  in  weighing  0.2  gram 
of  apatite  containing  40  %  P2O5,  what  will  be  the  apparent  per- 
centage ?  What  will  be  the  percentage  error  in  each  case  ? 

Answers:  65.01  %  Cu  ;  40.2  %  P2O5  ;  0.02  %  and  0.5  %. 

14.  If  the  dry  cupric  sulphide  from  0.82  gram  of  brass  loses 
0.1345  gram  on  ignition   in   hydrogen,  what  is  the  percentage  of 
copper  in  the  brass  ?     2  CuS  ==  Cu2S  +  S.      Answer  .  ^^  %< 

15.  If  1.5  grams  of  glass  yield  0.38  gram  KC1  -\-  NaCl,  from 
which  0.646  gram  2  KC1.  PtCl4  is  obtained,  what  is  the  percentage 
of  Na20  in  the  glass  ? 


1  6.  A  mixture  of  BaO  and  CaO  weighing  0.2438  gram  yields 
0.4876  grams  of  mixed  sulphates.  What  is  the  weight  of  each 
oxide  in  the  original  mixture  ? 

Answer:  0.1288  gram  CaO;  0.1150  gram  BaO. 

17.  Calculate  the   percentage  of  pure   Na2CO3  in   an   impure 
sample,  from  the  following  data  :    Crucible  -f-  SiO2  =  20.0697 
grams;  crucible  -f-  SiO2  -|-  Na2CO3  (impure)  =  20.3264  grams; 
crucible  -|-  SiO2  (excess)  -|-  Na2SiO3   (after  fusion)  =  20.2239 
grams.     Assume  the  reaction  to  be  Na2CO3  -f-  SiO2  =  Na2SiO3 
-j-  CO2.  Answer:  96.25  °/0. 

18.  A   sample  of  pyrite  weighing  0.5   gram  yields   1.6  grams 
BaSO4.     Calculate  the  percentage  FeS2  in  the  sample. 

Answer:  82.35  °/0. 

VOLUMETRIC   ANALYSIS. 

1.  How  much  crude  cream   of  tartar  should  be  taken  for  an 
analysis  in  order  that  the  number  of  cubic  centimeters  of  §  NaOH 
solution  required  to  react  with  it,  shall  represent  directly  the  per- 
centage of    KHC4H4O6  ?     How  much   oxalic   acid  in   order  that 
each  cubic  centimeter  of  T^  KMnO4  may  represent  i  %  C2H2O4. 
2  H2O  ?     KHC4H4O6  +  NaOH  =  KNaC4H4O6  +  H2O. 

Answer:  9.409  gram  KHC4H4O6  ;  0.6302  gram  C2H2O4.  2  H2O. 

2.  What  weight  of  potassium  ferrocyanide,  K4Fe(CN)6.  3  H2O, 
should   a  normal  solution  contain,  for  use   as  a  reducing  agent? 
10    K4Fe(CN)6.     (3     H2O)     +     2     KMnO4     +    8    H2SO4     = 
10  K8Fe(CN)6  +  6  K2SO4  +  2  MnSO4  +  8  H2O  +  (30  H2O). 


140 


PROBLEMS. 


3.  Calculate   the    percentage   of    carbon    dioxide   (CO2)    in    a 
sample   of  calcium  carbonate  from  the  following  data  : 

Total  volume  £  HC1  =  35  cc. ;  total  volume  ^  NaOH  ==  15  cc. ; 
weight  carbonate  =  i.oo  gram.  Answer :  35.20  °fc. 

4.  Calculate  the  weight  of   KHC2O4.  C2H2O4.  (2  H2O)  neces- 
sary for  a  liter  of  normal  solution,  (a)  as  a  standard  acid  solution 
(compare  note  7,  on  page  79),   (b)  as  a  reducing  agent.     KHC2O4. 
C2H2O4.  (2  H2O)    +    2   MnO2    +    3   H2SO4    =    2   MnSO4    + 
KHSO4  +  4  CO2  +  4  H2O  +  (2  H2O). 

Answer :  (a)  84.72  ;  (b)  63.54  grams. 

5.  Given  the  following  data,  calculate  the  percentage  purity  of 
the  oxalic  acid  : 

Standardization:  Weight  CaCO3  =  1.050  gram;  HC1  solution 
used  =  45  cc. ;  NaOH  solution  used  —  4.8  cc.  ;  i  cc.  NaOH 
solution  =  1.042  cc.  HC1  solution. 

Analysis:  Weight  oxalic  acid  =  1.500  gram;  NaOH  solution, 
used  =  42.5  cc. ;  HC1  solution  used  —  0.5  cc. 

Answer:  96.52   °/0. 

6.  Given  the  following  data,  calculate  the  percentage  purity  of 
the  cream  of  tartar  (KHC4H4O6)  : 

Weight  of  substance  =  2.500  grams  ;  NaOH  solution  used  = 
25.51  cc. ;  H2SO4  solution  used  =  0.5  cc. ;  i  cc.  H2SO4  solu- 
tion =  1.02  cc.  NaOH  solution;  i  cc.  NaOH  solution^  0.0255 
gram  CaCO3.  Answer:  95.90  °/0. 

7.  Solutions  of  alkali  carbonates  with  phenolphthalein  become 
colorless  as  soon  as  the  carbonate  has  changed  to  bicarbonate. 
Calculate  the  percentage  NaOH  in  a  sample  of  soda  ash  from  the 
following  data,  assuming  the  hydrate  to  be  neutralized  before  the 
carbonate    is    attacked  :    Weight   of    soda   ash  =  i   gram ;    HC1 
solution  is  £.     The  solution  becomes  colorless  when  25  cc.  HC1 
have    been    added,   but  requires  40  cc.  for   complete  neutraliza- 
tion, after  boiling  out  the  carbon   dioxide.     Na2COa  -f-  HC1  = 
HNaCO8  +  NaCl ;  NaHCO8  +  HC1  =  NaCl  +  CO2  +  H2O. 

Answer:  20.02  °/0. 

8.  If  10  cc.   of    a  sulphuric  acid  solution    yield  0.1220  gram 
BaSO4,  how  much  must  the  solution  be  diluted  for  an  exactly  ^ 
solution?  Answer:  1000  cc.  to  1044  cc. 

9.  If  i  cc.  of  a  potassium  bichromate  solution  will  oxidize  0.0066 
gram  iron,  to  what  volume  must  100  cc.  of  the  solution  be  diluted 
to  make  a  TJ^  solution  ?  Answer :  100  cc.  to  1178  cc. 


STOICHIOME  TRY.  141 

10.  Calculate  the  percentage  of  iron  (Fe)  in  a  sample  of  limon- 
ite  from  the  following  data  : 

Weight  of  limonite  =  0.55  gram;  K2Cr2O7  solution  used  = 
51.1  cc. ;  i  cc.  K2Cr2O7  solution  —  .0058  gram  Fe ;  FeSO4  solu- 
tion used  —  5  cc. ;  5  cc.  of  FeSO4  solution  contains  0.008  gram 
FeO.  Answer :  52.74  %. 

11.  A  sample  of  iron  wire  is  dissolved,  out  of  contact  with  air, 
in  30  cc.  HC1,  of  which    i   cc.  =  0.95  cc.  §  HC1.     The  iron  re- 
quires 40  cc.  T^  K2Cr2O7  for  oxidation.     What  excess  of  HC1  was 
used  over  that  required  for  solution?  Answer:  13.16  cc. 

12.  How  much   stannous   chloride  (SnCl2)   by  weight   will    it 
require  to  reduce  the  iron  from  0.5  gram  magnetite  (FeO.  Fe2O3), 
dissolved  out  of  contact  with  air?  Answer:  0.4093  gram. 

13.  How   many  cubic  centimeters  of   HC1  (sp.  gr.    1.12)   are 
required  to  dissolve  0.55  gram  limonite  (2  Fe2O3.  3  H2O),  assum- 
ing the  only  impurity  to  be  1.5  per  cent,  quartz? 

Answer:  2.37  cc. 

14.  If  0.75  gram  of  a  silicate  yields  0.4  gram  Fe2O3  -(-  A12O3, 
and  the  iron  present  requires  20  cc.  K2Cr2O7  solution  (i  cc.  = 
0.0784  gram  FeSO4(NH4)2SO4.  6  H2O),   calculate  the  percentage 
FeO  and  A12O3  in  the  sample. 

Answer:  38.37  %  FeO;   10.72  °/c  A12O3. 

15.  What  weight  of  iron  wire  containing  99.85  °]0  Fe  will  react 
with  the  chromium  from  0.5  gram  chromite  (FeO.  Cr2O3)? 

Answer  :  o  7504  gram. 

16.  If  i  cc.  of  a  potassium  permanganate  solution  will  oxidize 
0.0085  gram  Fe,  calculate  the  value  of  the  same  solution  in  terms 
of    (a)     KHC2O4.C2H2O4.  2  H2O  ;     (b)   KNO2  ;      (e)    Mn  ;     (<t) 
K4Fe(CN)6.  3  H2O. 

10  KHC2O4.  C2H2O4.  (2  H2O)   -f   8  KMnO4  -f-  17  H2SO4  = 

9  K2SO4   +   8  MnSO4   -f   40  CO2   +   32   H2O    +   (20  H2O); 

10  KNO2  +  4  KMnO4  +   n  H2SO4  =  7  K2SO4  +  4  MnSO4 
+    10  HNO3   -f-   6  H2O;    3  MnO   +   Mn2O7  =    5  MnO2,    and 
2  KMnO4  =  K2O.  Mn2O7 ;   compare  also  problem  2. 

Answer:  (a)  0.009640  gram;  (b)  0.006463  gram;  (t)  0.002502 
gram;  (d)  0.06414  gram. 

17.  Calculate  the  value  of  a  permanganate  solution,  of  which 
i  cc.  =  0.008  gram  Fe,  in  terms  of  MoO3.     Mn2O7  -f-  10  FeO 
=    5  Fe2O3   +    2  MnO  ;   7  Mn2O7  +  Mo24O37  =   24  MoO3   + 
14  MnO.  Answer:  0.007051  gram. 


142  PROBLEMS. 

18.  Given  the  following  data,  calculate  the  percentage  of  iron 
in  the  limonite  : 

Weight  of  limonite  —  0.55  gram  ;  KMnO4  solution  used  = 
30  cc  ;  i  cc.  KMnO4  solution  =  0.0084  gram  C2H2O4.  2  H2O. 

Answer  :  40.74  °/0. 

19.  The    calcium    oxalate    precipitate  from   0.5   gram   marble, 
when  treated  with  sulphuric  acid,  liberates  sufficient  oxalic  acid 
to  reduce  43  cc.  of  permanganate  solution  (i  cc.  =  0.01150  gram 
Fe).     Calculate  the  percentage  of  calcium  (Ca)  in  the  marble. 

Answer:  35.38  %. 

20.  If  i  cc.  KMnO4  solution  will  oxidize  0.008  gram  iron  (Fe), 
calculate  the  equivalent  of  the  same  solution  in  terms  of  hydrogen 
peroxide  (H2O2),  and  also  the  volume  of    oxygen  which   will  be 
evolved  by  each   cubic  centimeter  of  the   permanganate   solution 
during  the  reaction,  assuming  that  i  cc.  of  oxygen  weighs  0.00143 
gram   under   the    existing  conditions.     5  H2O2  -f-  2   KMnO4  -\- 
3  H2SO4  =  K2SO4.  +  2  MnSO4  +  5  O2  +  5  H2O. 

Answer:  0.002431  gram;   1.6  cc. 

21.  Given  the  following  data,  calculate  the  percentage  of  MnO2 
in  the  pyrolusite  : 

Weight  of  pyrolusite  =  0.48  gram  ;  weight  FeSO4.  (NH4)2SO4. 
6  H2O  =  4.3501  grams;  K2Cr2O7  solution  used  =  10  cc.  ;  i  cc. 
K2Cr207  solution  =  0.005  gram  Fe.  Ansmf  .  ^  ^  %_ 

22.  Given    the   following   data,    calculate    the    percentage    of 
MnO2  in  the  pyrolusite  : 

Weight  pyrolusite  =  0.48  gram  ;  weight  iodine  liberated  from 
Kl  =  1.296  grams. 

For  reactions,   see   note  6,   page   103.  Answer:   92.62    %. 

23.  If  i  cc.  iodine  solution  is  equivalent  in  oxidizing  power  to 
0.00149  gram  KBrO3,  to  what  volume  must  100  cc.  be  diluted  to 
make  a  „»„  solution  ? 


24.  Calculate  the  percentage  purity  of  the  sample  of  potas- 
sium bichromate  fro.m  the  following  data  : 

Weight  of  sample  =  0.1237  gram;  Na2S2O3  solution  used  = 
25  cc.  ;  i  cc.  Na2S2O3  solution  =  1.004  cc.  iodine  solution  ;  i  cc. 
iodine  solution  =  0.004975  gram  As2O3. 

K2Cr2O7  +  6  KI  -f  7  H2SO4  =  4  K2SO4  +  Cr2(SO4)3  + 
3  I2  +  7  H2O.  Answer  :  100  %. 


S  TOICHIOME  TR  Y.  1 43 

25.  Calculate  the  percentage  purity  of  a  sample  of  potassium 
iodate  (KIO3)  from  the  following  data  : 

Weight  of    sample  =  0.25  gram  ;    Na2S2O3  solution    used   = 
50  cc.  ;   i  cc.  Na2S2O3  solution  =  0.015  gram  I-    Answer :  84.34  °/0. 

26.  If  i  cc.  of  an  iodine  solution  has  the  same  oxidizing  power 
as  0.0034  gram  KIO3,  calculate  its  value  in   terms  of  antimony 
(Sb).  Answer:  0.005742  gram. 

27.  What  is  the  relation  of  each  of  the  following  solutions  to 
a  normal  solution  of  the  reagent  ? 

Na2S2O3.  5  H2O  :   i  cc.  =  0.001783  gram  KIO3. 

KMnO4  (acid  solution) :   i  cc.  =  0.006303  gram  C2H2O4.  2  H2O. 

As2O3  (as  a  reducing  agent)  :   i  cc.  =  0.01772  gram  Cl. 

As2O3  (as  an  acid)  :   i  cc.  =  0.01061  gram  Na2CO3. 

NaOH  :  i  cc.  =  0.1694  gram  KHC2O4.  C2H2O4.  2  H2O. 


144  CAPACITY  OF  LIPPED  BEAKERS. 


CAPACITY   OF   LIPPED    BEAKERS 
(REFERRED  TO  IN  THE  TEXT) 

WHEN    FILLED   TO    LOWER    EDGE    OF    LIP. 

No.  i 120  cc. 

NO.  2 210  CC. 

No.  3 310  cc. 

No.  4 540  cc. 

No.  5 700  cc. 

No.  6 930  cc. 

I 


ATOMIC  WEIGHTS. 

TABLE    OF   ATOMIC    WEIGHTS.* 
O  -  -  16.00. 


Aluminum 

Al 

27.11 

Mercury    .     . 

Hg 

200.0 

Antimony       .     . 

Sb 

120.43 

Molybdenum 

Mo 

95-98 

Argon   .... 

A 

? 

Neodyjnium   .     . 

Nd 

140.5 

Arsenic     .     .     . 

As 

75>°9 

Nickel  .... 

Ni 

58.69 

Barium       .     .     . 

Ba 

!37-43 

Nitrogen   .     .     . 

N 

14.04 

Bismuth     .     .     . 

Bi 

208.  1  1 

Osmium     .     . 

Os 

190.99 

Boron    .... 

B 

10.95 

Oxygen      .     . 

0 

16.0 

Bromine     .     .     . 

Br 

79-95 

Palladium       .     . 

Pd 

106.36 

Cadmium        .     . 

Cd 

111.93 

Phosphorus    .     . 

P 

31.02 

Caesium      .     .     . 

Cs 

132.89 

Platinum    .     .     . 

Pt 

194.89 

Calcium     .     . 

Ca 

40.08 

Potassium 

K 

39->i 

Carbon      .     .     . 

C 

12.01 

Praseodymium    . 

Pr 

I43-S 

Cerium      .     . 

Ce 

140.2 

Rhodium   .     .     . 

Rh 

103.01 

Chlorine    .     .     . 

Cl 

35-45 

Rubidium       .     . 

Rb 

85.43 

Chromium      .     . 

Cr 

52.14 

Ruthenium     .     . 

Ru 

101.68 

Cobalt        .     .     . 

Co 

58.93 

Samarium       .     . 

Sm 

150.0 

Columbium    .     . 

Cb 

94.0 

Scandium  .     . 

Sc 

44.0 

Copper      .     .     . 

Cu 

63  60 

Selenium   .     .     . 

Se 

79.0 

Erbium      .     .     . 

Er 

166.3 

Silicon       .     .     . 

Si 

28.40 

Fluorine     .     . 

Fl 

19.03 

Silver    .... 

Ag 

107.92 

Gadolinium    .     . 

Gd 

156.1 

Sodium       .     .     . 

Na 

23-°5 

Gallium      .     .     . 

Ga 

69.0 

Strontium       .     . 

Sr 

87.61 

Germanium     . 

Ge 

72.3 

Sulphur      .     . 

S 

32.07 

Glucinum        .     . 

Gl 

9.08 

Tantalum  .     . 

Ta 

182.6 

Gold      .... 

Au 

197.24 

Tellurium  .     . 

Te 

127.0 

Helium      .     .     . 

He 

? 

Thallium    .     .     . 

Tl 

204.15 

Hydrogen       .     . 

H 

1.0075 

Thorium    .     .     . 

Th 

232.63 

Indium       .     .     . 

In 

1  13  7 

Tin  

Sn 

I  TO   OC 

Iodine  .... 

I 

126.85 

Titanium   . 

Ti 

48.15 

Iridium 

Ir 

193.12 

Tungsten  .     . 

W 

184.84 

Iron       .... 

Fe 

56.02 

Uranium    .     .     . 

u 

239-59 

Lanthanum     .     . 

La 

138.6 

Vanadium       .     . 

V 

51.38- 

Lead     .... 

Pb 

206.92 

Ytterbium       .     . 

Yb 

173.0 

Lithium     .     .     . 

Li 

7-°3 

Yttrium      .     .     . 

Y 

88.95 

Magnesium     . 

Mg 

24.29 

Zinc       .... 

Zn 

65.41 

Manganese     . 

Mn 

54-99 

Zirconium       .     . 

Zr 

90.6 

*  F.  W.  Clarke,  /.  Am.  Chem.  Soc.t  18,  213. 


146  REAGENTS. 

STRENGTH    OF   REAGENTS.* 

GRAMS    PER    LITER. 

Ammonium  Molybdate  f  (of  MoO3)      . 68  grams. 

Ammonium  Oxalate,  (NH4)2C2O4.  H2O          ....     40  grams. 

Barium  Chloride,  BaCl2.  2  H2O 20  grams. 

Magnesium  Ammonium  Chloride  (of  MgCl2)     ...     64  grams. 

Mercuric  Chloride,  Hg<32        50  grams. 

Potassium  Hydroxide,  KOH  (sp.  gr.  1.27)     ....  480  grams. 

Potassium  Sulphocyanate,  KSCN 5  grams. 

Silver  Nitrate,  AgNO3 25  grams. 

Sodium   Hydroxide,  NaOH 100  grams. 

Sodium  Carbonate,  Na2CO8 150  grams. 

Sodium  Phosphate,  HNa2PO4.  12  H2O 100  grams. 

Stannous  chloride,  SnCl2,  made  by  saturating  hydrochloric  acid 
with  tin,  diluting  with  an  equal  volume  of  water,  and  adding 
a  slight  excess  of  acid  from  time  to  time.  A  strip  of  metallic 
tin  is  kept  in  the  bottle. 


Aqueous  ammonia,  sp.  gr.  0.96  contains  9.90%  NH3  by  weight, 
at  15°  C4 

Aqueous  ammonia,  sp.  gr.  0.90  contains  28.30%  NH3  by  weight, 
at  15°  C.t 

Ammonia  (0.96)  may  be  prepared  by  diluting  four  volumes  of  ammonia 
(0.90)  with  seven  volumes  of  water. 

Hydrochloric  acid,  sp.  gr.  1.12,  contains  23.8%  HC1  by  weight, 
at  15°  C.§ 

Hydrochloric  acid,  sp.  gr.  1.20,  contains  39.1%   HC1  by  weight, 
at  15°  C.§ 

Hydrochloric  acid  (1.12)  may  be  prepared  by  diluting  five  volumes  of 
hydrochloric  acid  (1.20)  with  four  volumes  of  water. 

Nitric  acid,  sp.  gr.   1.20,  contains  32.4%   HNO3  by  weight,  at 
i5°C.|| 

Nitric  acid,  sp.  gr.   1.42,  contains  69.8%   HNO3  by  weight,  at 

.5°  c.| 

Nitric  acid  (1.20)  may  be  prepared  by  diluting  two  volumes   of  nitric 
acid  (1.42)  with  three  volumes  of  water. 


*The  concentrations  given  in  this  table  are  those  upon  which  the  proce- 
dures in  the  foregoing  pages  are  based.  It  is  obvious,  however,  that  an  exact 
adherence  to  these  quantities  is  not  essential. 

t  This  solution  is  prepared  according  to  the  formula  of  Blair  and  Whitfield, 
/.  Am.  Chem.  Soc.,  17,  760. 

J  Lunge  and  Wiernik,  Ztschr.  angew.  Chem,  1889,  183. 

§  Lunge  and  Marschlewski,  Ztschr.  angeiv.  Chem,  1891,  133. 

||  Kolb.  Dingl.  pol.J.,  182,  233. 


1 
LOGARITHMS  OF  NUMBERS. 

Natural 
numbers. 

0 

I 

2 

3 

4 

5 

6 

7 

8 

9 

PROPORTIONAL  PARTS. 

I 

2 

3 

4 

5 

6 

7 

8 

9 

IO 

0000 

0043 

0086 

0128 

0170 

O2  I  2  0253 

0294 

°334 

0374 

4 

8 

12 

17 

21 

25 

29 

3337 

II 

0414 

0453 

0492 

053' 

0569 

0607  0645 

0682 

0719 

°755 

4 

8 

II 

ls 

19 

23 

26 

3034 

12 

0792 

0828 

oS6^ 

0899 

0934 

0969  1004 

103- 

1072 

1106 

3 

7 

io'i4 

1  7 

21 

24 

2831 

13 

"39 

i'73 

1206 

1239 

1271 

1303*335 

1367 

1  399*430 

3 

6 

1013 

1  6 

19 

23 

26  29 

14 

1461 

1492 

'523 

1553 

1584 

1614 

1644 

1673 

1703 

1732 

3 

6 

9 

12 

15 

18 

21 

24 

27 

IS 

1761 

1790 

1818 

1847 

1875 

1903 

I93l 

*959 

1987 

2014 

3 

6 

8 

II 

M 

17 

20 

22 

25 

16 

2041 

2068 

2095 

2122 

2148 

2I75!22OI 

2227 

2253 

2279 

3 

5 

8 

II 

'3 

16 

1  8 

21 

24 

'7 

2304 

2330 

2355 

2380 

2405 

2430  2455 

2480 

2504 

2529 

2 

5 

7 

IO 

12 

15 

J7 

20 

22 

18 

2553 

2577 

2601 

2625 

2648 

2672 

2695 

2718 

2742 

2765 

2 

5 

7 

9 

12 

M 

16 

19 

21 

19 

2788 

2810 

2833 

2856 

2878 

29OO 

2923 

2945 

2967 

2989 

2 

4 

7 

9 

II 

^3 

16 

iS 

20 

20 

3010 

3032 

3054 

3075 

3096 

3II8 

3i39 

3160 

3181 

3201 

2 

4 

6 

8 

II 

*3 

»5 

17 

19 

21 

3222 

3243 

3263 

3284 

33°4 

3324 

3345 

3365 

3385 

3404 

2 

4 

6 

8 

10 

12 

M 

16 

18 

22 

3424 

3444 

3464 

3483 

3502 

3522 

354i 

3560 

3579 

3598 

2 

4 

6 

8 

10 

12 

M 

I5i7 

23 

3617 

3636 

3655 

3674 

3692 

3711 

3729 

3747 

3766 

3784 

2 

4 

6 

7 

9 

II 

J3 

15 

17 

24 

3802 

3820 

3838 

3856 

3874 

3892 

3909 

3927 

3945 

3962 

2 

4 

5 

7 

9 

II 

12 

M 

16 

25 

3979 

3997 

4014 

4031 

4048 

4065 

4082 

4099 

4116 

4133 

2 

3 

5 

7 

9 

10 

12 

M 

T5 

26 

415° 

4166 

4183 

4200 

4216 

4232 

4249 

4265 

4281  4298 

2 

3 

5  7 

8 

IO 

II 

iji5 

27 

43M 

433° 

4346 

4362 

4378 

4393 

4409 

4425 

4440  4456 

2 

3 

5  6 

8 

9 

II 

'3  X4 

28 

4472 

4487 

4502 

45'8 

4533 

4548 

4564 

4579 

459414609 

2 

3 

5  6 

8 

9 

II 

12  14 

29 

4624 

4639 

4654 

4669 

4683 

4698 

4713 

4728 

4742 

4757 

I 

3 

4 

6 

7 

9 

10 

12 

13 

30 

4771 

4786 

4800 

4814 

4829 

4843 

4857 

4871 

4886 

4900 

I 

3 

4 

6 

7 

9 

10 

I  I 

13 

31 

4914 

4928 

4942 

4955 

4969 

4983 

4997 

5011 

5024 

5038 

I 

3 

4 

6 

7 

8 

10 

II 

12 

32 

5051  5065 

5079 

5092 

5I05 

5IJ9 

5132 

5M5 

5i59 

5172 

I 

3 

4 

5 

7 

8 

9 

I  I 

12 

33 

5185 

5i9S 

5211 

5224 

5237 

525° 

5263 

5276 

5289  5302 

I 

3 

4 

5 

6 

8 

9 

10 

12 

34 

53r5 

5328 

5340 

5353 

5366 

5378 

539i 

5403 

54i6 

5428 

I 

3 

4 

5 

6 

8 

9 

IO 

II 

35 

544i 

5453 

5465 

5478 

5490 

5502 

55J4 

5527 

5539 

555i 

I 

2 

4 

5 

6 

7 

9 

10 

II 

36 

5563 

5575 

5587 

5599 

5611 

5623 

5635 

5647 

5658 

5670 

I 

2    4 

5 

6 

7 

8 

10 

I  1 

37 

5682 

5694 

5705 

57i7 

5729 

5740 

5752 

5763 

5775 

5786 

I 

**   o 

5 

6 

7 

8 

9 

10 

38 

5798 

5809 

5821 

5832 

5843 

5855 

5866 

5877 

5888 

5899 

I 

2 

3 

5 

6 

7 

8 

9  10 

39 

5911 

5922 

5933 

5944 

5955 

5966 

5977 

5988 

5999 

6010 

I 

2 

3 

4 

5 

7 

8 

9  10 

40 

6021 

6031  6042 

6053 

6064 

6075 

6085 

6096 

6107 

6117 

2 

3 

4 

5 

6 

8 

910 

4i 

6128 

6138  6149 

6160 

6170 

6180 

6191 

6201 

6212 

6222 

2 

3 

4 

5 

6 

7 

8 

9 

42 
43 

6232 
6335 

6243  6253 
6345  6355 

6263 
6365 

6274 
6375 

62841629463046314 
6385  .6395  6405  64  1  5 

6325 
6425 

2 

3 
3 

4 
4 

5 

5 

6 
6 

7 
7 

8 
8 

9 

9 

44 

6435 

6444 

6454 

6464 

6474 

6484 

6493 

6503 

65'3 

6522 

2 

3 

4 

5 

f,  7 

8  9 

45 

6532 

6542 

6551 

6561 

6571 

6580 

6590 

6599 

6609 

6618 

2 

3 

4 

5 

6 

7 

8 

9 

46 

6628 

6637 

6646 

6656 

6665 

6675 

6684 

6693 

6702 

6712 

i 

3 

4 

5 

6 

7 

7 

8 

47 

6721 

6730 

6739 

6749 

6758 

6767 

6776 

6785 

6794 

6803 

2 

3 

4 

5 

5 

6 

7 

8 

48 

6812 

6821 

6830 

6839 

6848 

6857 

6866 

6875 

6884 

6893 

2 

3 

4 

4 

5 

6 

7 

8 

49 

6902 

6911 

6920 

6928 

6937 

6946 

6955 

6964 

6972 

6981 

2 

3 

4 

4 

5 

6 

7 

8 

5° 

6990 

6998 

7007 

7016 

7024 

7033 

7042 

7050 

7059 

7067 

2 

3 

3 

4 

5 

6 

7 

8 

51 

7076 

7084 

7093 

7101 

7110 

7118 

7126 

7M5 

7M3 

7152 

2 

3 

3 

4 

5 

6 

7 

8 

52 

7160 

7168 

7177 

7185 

7i93 

7202 

7210 

7218 

7226 

"235 

2    2 

3 

4 

5 

6 

7 

7 

53 

7243 

7251 

7259 

7267 

7275 

7284 

7292 

7300 

7308 

7316 

2    2 

3 

4 

5 

6 

6 

7 

54 

7324 

7332 

7340  7348 

7356 

7364 

7372 

738o 

7388 

7396 

2    2 

3 

4 

5 

6 

6 

7 

LOGARITHMS  OF  NUMBERS. 

Natural 
numbers. 

O 

I 

2 

3 

4 

5 

6 

7 

8 

9 

PROPORTIONAL  PARTS. 

I 

2 

3 

4 

5 

6 

7 

8 

9 

55 

7404 

7412 

74197427 

7435 

7443 

7451 

7459 

7466 

7474 

, 

2 

2 

3 

4 

5 

5 

6 

7 

56 

7482 

7490 

7497  75°5 

75r3 

7520 

7528 

7536 

7543 

7551 

I 

2 

2 

3 

4 

5 

5 

6 

7 

57 

7  $59 

7566 

75747582 

7589 

7597  7604 

7612 

7619 

7627 

I 

2 

2 

3 

4 

5 

5 

6 

7 

58 

7634 

7642 

7649 

7657 

7664 

7672  7679  7686 

7694 

7701 

I 

2 

3 

4 

4 

5 

6 

7 

59 

7709 

7716 

7723 

7731 

7738 

77457752776o 

7767 

7774 

I 

2 

3 

4 

4 

5 

6 

7 

60 

7782 

7789 

7796 

7803 

7810 

78187825  7832 

783917846 

I 

2 

3 

4 

4 

5 

6 

6 

6r 

7853 

7860 

7868 

7875 

7882 

7889:78967903 

7910 

7917 

I 

2 

3 

4 

4 

5 

6 

6 

62 

7924 

793i 

7938 

7945 

7952 

7959  7966  7973 

798o 

7987 

I 

2 

3 

3 

4 

5 

6 

6 

63 

7993 

8000 

8007 

8014 

8021 

802818035 

8041 

8048 

8055 

I 

2 

3 

3 

4 

5 

5 

6 

64 

8062 

8069 

8075 

8082 

8089 

8096 

8102 

8109 

8116 

8122 

I 

2 

3 

3 

4  5 

5 

6 

65 

8129 

8136 

8142 

8149 

8156 

8162 

8169 

81768182 

8189 

I 

I 

2 

3 

3 

4  5 

5 

6 

66 

8195  8202 

8209 

8215 

8222 

8228 

8235 

8241  ^24818254 

I 

I 

2 

3 

3 

4   S 

5 

6 

67 

8261 

8267 

8274 

8280 

8287 

8293 

8299 

830683128319 

I 

2 

3 

3 

4  5  5 

6 

68 

8325 

833i 

8338 

8344 

835i 

8357 

8363 

837083768382 

I 

*y 

3 

3 

445 

6 

69 

8388 

8395 

8401 

8407 

8414 

8420 

8426 

8432  8439  8445 

I 

2 

2 

3 

4  4 

5 

6 

70 

8451 

8457 

8463 

8470 

8476 

84828488 

8494  8500  8506 

I 

2 

2 

3 

4  4 

5 

6 

7i 

8513 

8519 

8525 

853i 

8537 

85438549 

855585618567 

I 

2 

2 

3 

4  4  5 

5 

72 

8573 

8579 

8585 

859i 

8597 

8603 

8609 

861518621,8627 

I 

2 

2 

3 

44  5 

5 

73 

8633 

8639 

8645 

8651 

8657 

8663 

8669 

86758681 

8686 

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2 

2 

3 

4.4  5 

5 

74 

8692 

8698 

8704 

8710 

8716 

8722 

8727 

8733  8739 

8745 

I 

2 

2 

3 

4  4  5 

5 

75 

8751 

8756 

8762 

8768 

8774 

8779 

8785 

8791 

8797 

8802 

I 

2 

2 

3 

3  4  5 

5 

76 

8808 

8814 

88208825 

8831 

8837 

8842 

8848 

8854 

8859 

I 

2 

2 

3 

34  5 

5 

77 

88658871 

8876 

8882  8887 

88938899 

8904 

89108915 

I 

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2 

3 

344 

5 

78 

8921  8927 

8932 

8938 

8943 

8949^8954 

8960 

89658971 

I 

2 

2 

3  3 

4  4 

5 

79 

8976 

8982 

8987 

8993 

8998 

9004 

9009 

9°!  5 

9020  9026 

l 

2 

2 

3 

3 

4  4 

5 

80 

9031 

9036 

9042 

9047 

9053 

9058 

9063 

9069 

9074  9079 

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2 

2 

3 

3 

4  4 

5 

81 

9085  9090 

9096 

9101 

9106 

9112 

9117 

9122 

91289133 

I 

2 

2 

3  3 

4 

4 

5 

82 

9*38  9M3 

9149 

9r5-l 

9159 

9^5 

9170 

9*75 

91809186 

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2 

2 

3 

3 

4 

4  5 

83 

9191 

9196 

9201 

9206 

9212 

9217 

9222 

9227 

9232 

9238 

I 

2 

2 

3 

3 

4 

4 

5 

84 

9243 

9248 

9253 

9258 

9263 

9269 

9274 

9279 

9284 

9289 

I 

2 

2 

3 

3 

4 

4 

5 

85 

9294 

9299 

93°4 

93°9 

93*5 

9320 

9325 

933° 

9335 

9340 

I 

2 

^ 

3 

3 

4 

4  5 

86 

9345 

9350 

9355 

9360 

9365 

9370 

9375 

9380 

9385 

9390 

I 

2 

2 

3 

3 

4 

4  5 

87 

9395 

9400 

9405 

9410 

94i5 

9420 

9425 

943° 

9435 

9440 

0 

2 

2 

3 

3 

4  4 

88 

9445 

9450 

9455 

9460 

9465 

9469 

9474 

9479 

9484 

9489 

0 

2 

s-> 

3 

3 

4  4 

89 

9494 

9499 

9504 

95°9 

95'3 

95i8 

9523 

9528 

9533 

9538 

O 

2 

2 

3 

3 

4  4 

90 

9542 

9547 

9552 

9557 

9562 

9566 

9571 

9576 

958  1  '9586 

0 

2 

2 

3 

3 

4  4 

9i 
92 

9590 
9638 

9595  9600 
96439647 

9605 
9652 

9609 
9657 

9614 
9661 

9619 
9666 

9624 
9671 

9628  9633 
96759680 

0 
0 

0 

2 

2 

3 
3 

3 
3 

4  4 
4  4 

93 

9685 

9689 

9694 

9699 

9703 

97089713 

9717 

9722 

9727 

0 

2 

2 

3 

3 

4  4 

94 

973i 

9736 

974i 

9745 

975° 

97549759 

9763 

9768 

9773 

0 

2 

2 

3 

3 

4  4 

95 

9777 

9782 

9786 

9791 

9795 

98009805 

9809 

9814 

9818 

0 

, 

2 

3 

3 

4  4 

96 

9823 

9827 

9832 

9836 

9841 

9845  9850 

9854 

9859 

9863 

O 

2 

J 

3 

3 

4  4 

97 

9868 

9872 

9877 

9881 

9886 

9890  9894 

9899 

9903 

9908 

O 

2 

2 

3 

3 

4  4 

98 

9912 

9917 

9921 

9926 

993° 

9934  9939 

9943 

994s 

9952 

O   I 

2 

2 

3 

3 

4  4 

99 

9956 

9961 

9965 

9969 

9974 

9978  9983 

9987 

999' 

9996 

0    I 

2 

2 

3 

3 

3]  4 

ANTILOGARITHMS. 

J 
5* 

O   I 

2 

3 

4 

5 

6 

7 

8 

9 

PROPORTIONAL  PARTS. 

I 

2 

3 

4 

5 

6 

7 

8 

9 

.00 

1000 

1  002  1005 

1007 

1009 

IOI2 

1014 

1016 

IOI9  IO2I 

O 

0 

i 

I 

2   2 

2 

.01 

1023 

1026  1028 

1030 

1033 

1035 

1038 

1040 

IO42  IO45 

O 

O 

i 

I 

2   2 

2 

.02 

1047 

1050  1052 

1054 

1057 

1059 

1062 

1064 

1067 

1069 

0 

O 

i 

I 

2   2 

2 

.03 

1072 

1074  1076 

1079 

1081 

1084 

1086 

1089 

I09T 

1094 

0 

0 

i 

i 

2 

2 

2 

.04 

1096 

1099 

1  102 

1104 

1107 

1109 

1  1  12 

1114 

III7 

1119 

O 

1 

i 

2 

2 

2 

2 

•  O  C 

1122 

1125 

1127 

"3° 

1132 

"35 

1138 

1140 

"43 

1146 

0 

i 

2 

2 

2 

2 

.06 

1148 

"51 

"53 

(156 

"59 

1161 

1164 

1167 

1169 

1172 

O 

i 

2 

2 

2 

2 

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"75 

1178 

1180 

"83 

1186 

1189 

II9I 

"94 

"97 

"99 

O 

i 

2 

2 

2 

2 

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1  202 

1205 

1208 

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1213 

1216 

1219 

1222 

1225 

1227 

0 

i 

2 

2 

2 

3 

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1230 

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1236 

12391242 

1245 

1247 

I25O 

1253 

1256 

0 

i 

2 

2 

2 

3 

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1259 

1262 

1265 

12681271 

1274  1276 

1279 

1282 

1285 

O 

i 

2 

2   2 

3 

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1288 

1291 

1294 

1297  1300 

1303  1306 

1309 

1312 

1315 

0 

2 

2 

2   2 

3 

.12 

1318 

1321 

1324 

132711330 

1334  1337 

1340 

1343 

1346 

0 

2 

2 

2   2 

3 

•'3 

1349 

1352 

1355 

r358  1361 

1365  1368 

1371 

1374 

'377 

0 

2 

2 

2 

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3 

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1380 

1384 

1387 

1390 

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1396 

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1403 

1406 

1409 

0 

2 

2 

2  3 

3 

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.16 

1413 

1445 
M79 

1416 
1449 
1483 

1419 
1452 
1486 

1422 
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1426 
1459 
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1429 
1462 
1496 

1432 
1466 
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M35 
1469 

1503 

M39 
1472 

1442 
1476 
1510 

0 
0 
0 

2 
2 
2 

2 
2 

2 

3 

*3 

3 
3 
3 

'.18 
.19 

1514 
1549 

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1524 
1560 

1528 
1563 

1531 
1567 

1535 
1570 

1538 
1574 

1542 
1578 

1545 
1581 

0 

0 

2 

2 

2 
2 

2  i  3 

3  3 

3 
3 

.20 

1585 

i  589  i  592 

1596 

1600 

1603 

1607 

1611 

1614 

1618 

0 

i 

2 

2 

3 

3 

3 

.21 

1622 

1626  1629 

1633  1637 

1641 

1644 

1648 

1652 

1656 

O 

2 

2 

2   3 

3 

3 

.22 

1660 

1663  1667 

1671 

1675 

1679 

l683 

1687 

1690 

1694 

O 

2 

2 

3 

3 

3 

•23 

1698 

170211706 

1710 

1714 

1718 

1722 

1726 

173° 

1734 

0 

2 

2 

2 

3 

3 

4 

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1738 

174^1746 

i75OI754 

1758 

I762 

1766 

1770 

1774 

0 

2 

2 

2  3 

3 

4 

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1778 

1782  1786 

i79I|I795 

1799 

I803 

1807 

1811 

1816 

O 

2 

2 

2  3 

3 

4 

.26 

1820 

1824 

1828 

1832  1837 

1841 

1845 

1849 

1854 

1858 

0 

2 

2  3-  3 

3 

4 

•27 

1862 

1866 

1871 

,875  1879 

1884 

1888 

1892 

1897 

1901 

0 

2 

2 

3 

3 

3 

4 

.28 

1905 

1910 

1914 

19191923 

1928 

1932 

1936 

1941 

1945 

0 

2 

2 

3 

3 

4 

4 

.29 

195° 

J954 

J959 

1963  1968 

1972 

1977 

1982 

1986 

1991 

0 

2 

2 

3 

3 

4 

4 

•3° 

J995 

2000 

2004 

2009  2014 

2018 

2023 

2028 

2032 

2037 

0 

2 

2 

3 

3  4 

4 

•31 

2042 

2046 

2051 

2056  2061 

2065 

2070 

2075 

2080 

2084 

0 

2 

2 

334 

4 

•32 

2089 

2094 

2099 

2104  2109 

2113 

21  18 

2123 

212812133 

0 

2 

2 

3  3  4 

4 

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2138 

2143 

2148 

2153,2158 

2163 

2168 

2I73 

2178(2183 

0 

2 

2 

3  3  4 

4 

•34 

2188 

2193 

2198 

22032208 

2213 

2218 

2223 

2228 

2234 

1 

2 

2 

3 

3  4  4 

5 

•35 

2239 

2244 

2249 

2254  2259 

2265 

2270 

2275 

2280  2286 

2   2 

3 

344 

5 

•36 

•37 

2291 
2344 

2296 
2350 

2301 

2355 

230712312 
2360  2366 

2317 
2371 

2323 

2377 

2328 
2382 

23332339 
2388  2393 

2   2 

2   2 

3  !  3  1  4  4 

3i3  4  4 

5 
5 

•38 
•39 

2399 

2455 

2404 
2460 

2410 
2466 

2415 

2472 

2421 

2477 

2427 
2483 

2432 
2489 

2438 
2495 

2443  2449 
250012506 

2   2 

2   2 

3 
3 

3  '  4  !  4 
3  4  5 

5 
5 

.40 

2512 

2518 

2523 

25292535 

2541 

2547 

2553 

2559:2564 

2 

2 

3445 

5 

.41 

2570 

2576 

2582 

2588  2594 

2600 

2006 

2612 

2618  2624 

2 

2 

3  4415 

5 

.42 

2630 

2636 

2642  2649  2655 

2661 

2667 

2673 

2679  2685 

2 

2 

3 

4  4  5 

6 

•43 

2692 

2698 

2704 

2710  2716 

2723 

2729 

2735 

274212748 

2 

3 

3 

4  |4  i  5 

6 

•44 

2754 

276^2767 

2773  2780 

2786 

2793 

2799 

2805  2812 

2 

3 

3 

4  4  5 

6 

•45 

2818 

28252831 

283812844 

2851 

2858 

28642871  2877 

2 

3 

3 

4 

5  5 

6 

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2884 

2891 

2897 

290412911 

2917 

2924 

2931  293812944 

1  2 

3 

3 

4 

5  5 

6 

•47 
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2951 
3020 
3090 

2958 
3027 
3097 

2965 
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2972 

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2979 
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2985 
3055 
3126 

2992 
3062 
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2999 
3069 
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30063013 
3076  3083 

2 
2 

2 

3 

3 
3 

3 
4 
4 

4 
4 
4 

5  5 
5  6 
5  ^ 

6 
6 
6 

1 

ANTILOGARITHMS. 

i.J 

O 

I 

2 

3 

4 

5 

6 

7 

8 

9 

PROPORTIONAL  PARTS. 

I 

2 

3 

4 

5 

6 

7 

8 

9 

.50 

•51 

3162 
3236 

3170 
3243 

3177 
3251 

1 
31843192 
3258  3266 

3T99 

3273 

3206 
3281 

3214 
3289 

3221 
3296 

3228 
3304 

I 
I 

2 

2 
2 

3 
3 

4 
4 

4 
5 

c 

i 

6 
6 

7 
7 

•52 

3311 

33*9 

3327 

33343342 

335° 

3357 

5365 

337  3  |338i 

I 

2 

2 

3 

4 

5 

e 

6 

7 

•53 

3388 

3396 

3404 

3412 

3420 

3428 

3436 

3443 

345  1  134  59 

I 

2 

2 

3 

4 

5 

6 

6 

7 

•54 

3467 

3475 

3483 

3491 

3499 

35°8 

35i6 

3524 

3532|3540 

I 

2 

2 

3 

4 

5 

6 

6 

7 

•55 

3548 

3556 

3565 

3573 

358i 

3589 

3597 

3606 

3614  3622 

I 

2 

2 

3 

4 

5 

6 

7 

7 

•56 

3639 

3648 

3656 

3664 

3673 

3681 

3690 

3698  3707 

I 

2 

3 

3 

4 

5 

6 

7 

8 

•59 

3715 
3802 
3890 

3724 
3811 

3899 

3733 
3819 
3908 

374i 
3828 

3917 

375° 
3837 
3926 

3758 
3846 

393° 

3767 
3855 
3945 

3776 
3864 
3954 

3784  3793 
3873  3882 

I 
I 

2 
2 
2 

3 
3 
3 

3 
4 
4 

4 
4 

5 

5 
5 
5 

6 
6 
6 

7 

7 
7 

8 
8 
8 

.60 

398i 

3990 

3999 

4009 

4018 

4027 

4036 

4046 

4055 

4064 

I 

2 

3 

4 

5 

6 

6 

7 

8 

.61 

4074 

4083 

4093 

4102 

4111 

4121 

413° 

4140 

4150 

4159 

I 

2 

3 

4 

5 

6 

7 

8 

9 

.62 

4169 

4178 

4188 

4198 

4207 

4217 

4227 

4236 

4246 

4256 

I 

2 

3 

4 

5 

6 

7 

8 

9 

i 

4266 
4365 

4276 
4375 

4285 
4385 

4295 
4395 

4305 
4406 

4416 

4325 
4426 

4335 
4436 

4345  4355 
444614457 

I 
I 

2 
2 

3 

3 

4 
4 

5 
5 

6 
6 

7 

7 

8 
8 

9 
9 

2 

4467 

457  1 

4477 
4581 

4487 
4592 

4498 
4603 

4508 
4613 

45T9 
4624 

4529 
4634 

4539 
4645 

4550 
4656 

4560 
4667 

I 
I 

2 

2 

3 

3 

4 
4 

5 

5 

6 
6 

7 

7 

8 
9 

9 

IO 

.67 

4677 

4088 

4699 

4710 

4721 

4732 

4742 

4753 

4764 

4775 

I 

2 

3 

4 

5 

7 

8 

9 

IO 

68 

4786 

4797 

4808 

4819 

4831 

4842 

4853 

4875 

4887 

I 

2 

3 

4 

6 

7 

8 

9 

IO 

'.69 

4898 

4909 

4920 

4932 

4943 

4955 

4966 

4977 

4989 

5000 

I 

2 

3 

5 

6 

7 

8 

9 

IO 

.70 

5012 

5023 

5035 

5°47 

5058 

5070 

5082 

5093 

5105 

5117 

I 

2 

4 

5 

6 

7 

8 

9 

II 

5129 

5152 

5164 

5^6 

5188 

5200 

5212 

5224 

5236 

I 

2 

4 

5 

6 

7 

8 

10 

II 

.72 

5248 

5260 

5272 

5284 

5297 

53°9 

5321 

5333 

5346 

5358 

I 

2 

4 

5 

6 

7 

9 

IO 

II 

•73 

5370 

5383 

5395 

5408 

5420 

5433 

5445 

5458 

5470 

5483 

I 

3 

4 

5 

6 

8 

9 

IO 

II 

•74 

5495 

5508 

5534 

5546 

5559 

5572 

5585 

5598 

5610 

I 

3 

4 

5 

6 

8 

9 

10 

12 

•75 

5623 

5636 

5649 

5662 

5675 

5689 

5702 

57  x  5 

5728 

5741 

I 

3 

4 

5 

7 

8 

9 

IO 

12 

.76 

5754 

5768 

578i 

5794 

5808 

5821 

5834 

5848 

5861 

5875 

I 

3 

4 

5 

7 

8 

9 

II 

12 

•77 

5888 

5902 

5929 

5943 

5957 

5970 

5984 

5998 

6012 

I 

3 

4 

7 

8 

10 

II 

12 

.78 

6026 

6039 

6053 

6081 

6095 

6109 

6124 

6138 

6152 

I 

3 

4 

6 

7 

8 

10 

II 

13 

•79 

6166 

6180 

6194 

6209 

6223 

6237 

6252 

6266 

6281 

6295 

I 

3 

4 

6 

7 

9  10 

II 

*3 

.80 

6310 

6324 

6339 

6353 

6368 

6383 

6397 

6412 

6427 

6442 

I 

3 

4 

6 

7 

9  10 

12 

I3 

.81 
.82 

6457 
6607 

6471 
6622 

6486 
6637 

6501 
665. 

6516 
6668 

653' 
6683 

6546 
6699 

6561 
6714 

6577 
6730 

6592 
6745 

2 

2 

3 
3 

5 

5 

6 
6 

8 
8 

9  " 
9  ii 

12 
12 

•4 

•83 

6761 

6776 

6792 

6868 

6823 

6839 

6855 

6871 

6887 

6902 

2 

3  5  6 

8 

9  II 

13 

H 

.84 

6918 

6934 

6950 

6966 

6982 

6998 

7015 

7031 

7047 

7063 

2 

3  5  6 

8 

10  II 

13 

'5 

•85 

7079 

7096 

7112 

7129 

7M5 

7161 

7178 

7i94 

7211 

7228 

2 

3 

5 

7 

8 

IO  12 

I3 

15 

.86 

7244 

7261 

7278 

7295 

73" 

7328 

7345 

7362 

7379 

7396 

2 

3 

5 

7 

8 

IO:  12 

'3 

'5 

.87 
.88 
•89 

7413 
7586 
7762 

743° 
7603 
7780 

7447 
7621 

7798 

7464 
7638 
7816 

7482 
7656 
7834 

7499 
7674 
7852 

75l6 
7691 
7870 

7534 
7709 
7889 

755' 
7727 
7907 

7568 
7745 
7925 

2 
2 
2 

3 
4 
4 

5 
5 
5 

7 
7 
7 

9 
9 
9 

10  12 

II  12 
"  13 

14 
14 

16 
16 
16 

.90 

7943 

7962 

7980 

7998 

8017 

8035 

8054 

8072 

8091 

8110 

2 

4 

6 

7 

9 

"i  13 

15 

17 

.91 

8128 

8147 

8166 

8185 

8204 

8222 

8241 

8260 

8279 

8299 

2  4 

6 

8 

9 

II 

'3 

15 

17 

.92 
•93 
•94 

8318 
8511 
8710 

8337 
8531 
8730 

8356 
8750 

8375 
8570 
8770 

8395 
8590 
8790 

8414 
8610 
8810 

f433 
8630 

8831 

8453 
8(350 
8851 

8472  8492 
867018690 
887218892 

2 
2 
2 

4 
4 
4 

6 
6 
6 

8 
8 
8 

10 
10 
10 

12 
12 
12 

>4 
M 
M 

15 
16 

£.0)00 

•95 

8913 

8933 

8954 

8974 

8995 

9016 

9036 

9057 

90789099 

2 

4  6 

8 

IO 

12  15 

17 

19 

.96 

9120 

9141 

9162 

9183 

9204 

9226 

9247 

9268 

9290931  r 

2 

4  6 

8 

1  1 

I  7 

15 

17 

19 

•97 

9333 

9354 

9376 

9397 

9419 

9441 

9462  9484 

95°6;95-8 

2 

4  7  j  9 

II 

13 

J5 

20 

.98 

955° 

9572 

9594 

9616 

9638 

9661 

9683 

9705 

9727 

975° 

2 

4 

7  i  9 

II 

'3 

16 

18 

2O 

•99 

9772 

9795 

9817 

9840 

9863 

9886 

9908 

993  : 

9954 

9977 

2  5 

7 

9 

II 

16 

18 

2O 

INDEX. 


Acidimetry 74 

Acid  solutions,  normal    ....  65 

standard      ...  74 
Acids,    weak,    actions    of    other 

acids  on 126 

Acids,  weak,  action  of  salts  on      -125 

Accuracy  demanded 20 

Alkalimetry 74 

Alkali  solutions,  normal      ...  65 

standard  .     .     .  74 

Ammonium  nitrate,  acid      ...  36 

Antimony,  determination  of    .     .  108 

Apatite,  analysis  of 36 

Asbestos  filters      .     .     .     .14,  27,  28 

Atomic  weights,  table  of      ...  145 

Balances,  use  and  care  of    .     .     .  16 

Bichromate  process  for  iron     .     .  85 

Brass,  analysis  of 50 

Burette,  description  of    ....  66 

calibrat'on  of     ....  68 

cleaning  of 70 

reading  of     .     .     .     .    69,  70 

Calcium,  determination  of  .     .     .  41 

Calibration,  definition  of     ...  67 

of  burettes  .     .     .     .  68 

of  flasks 71 

Carbon  dioxide,  determination  of,  45 

Chlorimetry in 

Chlorine,   gravimetric  determina- 
tion   21,  27 

Chlorine,    volumetric    determina- 
tion    in 

Chrome  iron  ore,  analysis  of    .     .  92 

Colloidal  precipitates      .     .     .     .  12 

Copper,  determination  of    ...  52 

Crucibles,  use  of 9 

Crystalline  precipitates  ....  12 


Desiccators  .     .     .     . 
Direct  methods      .     . 
Dissociation,  degree  of 
Dolomite,  analysis  of 


9 
64 


Economy  of  time 19 

Electrolytes,  interaction  of .     .     .123 
Electrolytic    dissociation,    theory 

of u6 

End-point,  definition  of  ....  63 

Equilibrium,  chemical     ....  119 

Evaporation  of  liquids    .      .     .     .  10 

Feldspar,  analysis  of 60 

Ferrous-ammonium  sulphate,  an- 
alysis of 29 

Filters,  how  fitted 13 

Filtration 13 

Flasks,  calibration  of      ....  71 

graduated 67 

Funnels 13 

Gelatinous  precipitates    .     .     .     .  15 

General  directions 8 

General  directions  for  volumetric 

analysis 72 

Gooch  filter 14,  27,  28 

Gravimetric  analysis 21 

definition  of,  7 


Hardened  filters 
Hydrolysis  .  . 
Hydrolytic  action 


.     ...      14 

.     ...   128 
,.  128,  130,  131 


Ignition  of  precipitates  .     .     .     .      i  t> 

Indicators,  definition  of  ....     63 

for  acidimetry    .     .  7  5,  1 29 

Indirect  methods 64 

Integrity  demanded 20 

lodimetry 104 

Ions,  definition      ....      116,  117 

color  of 119 

complex    .     .     .     .     .     123,  127 

lonization 118 

Iron,  gravimetric  determination  of,     29 
volumetric    determination 
of 90,  100 


Laumontite,  analysis  of  . 
Lead,  determination  of    . 


INDEX. 


'53 


Limonite,   determination    of   iron 

in 90,  100 

Liquids,  evaporation  of    .     .     .     .  10 

Logarithms 147 

Magnesium,  determination  of  .     .  43 

Mass  action,  law  of 120 

Mohr  liter  .                                       .  68 


Neatness 8 

Normal  solutions 64 

advantages  of     ....  66 

preparation  of     ....  76 

Oxalic  acid,  strength  of  ....  82 
Oxidation  processes  .  .  .  .  63,  84 
Oxidizing  solutions,  normal  .  .  65 

Permanganate  process  for  iron  .  94 
Phenolphthalein,  reaction  of  with 

carbonates Sr,  131 

Phenolphthalein,  reaction  of  with 

ammonia     ' 73,  130 

Phosphoric  anhydride,  determina- 
tion of     36 

Pipette,  description  of  ....  67 
Platinum  crucibles,  care  of  .  .  .  10 

Precipitates,  colloidal 12 

crystalline  ....  12 
re-solution  of  .  .  .126 
separation  from  filter,  24 

Precipitation 12,  121 

Precipitation  methods       .     .       63,  113 

Problems 137 

Pyrolusite,  oxidizing  power  of.     .   102 

Reagents 1 1 

Reducing  solution,  normal  ...  65 
Reductor,  Jones 96 


Saturation  methods  .  .  .  .  63,  74 
Silica,  determination  of  .  .  .  57,  60 
Silver,  determination  of  .  .  .  .114 
Soda  ash,  alkaline  strength  of  .  .  80 
Sodium  carbonate,reaction  of  with 

mineral  acids 130 

Sodium  chloride,  analysis  of     .     .     22 
preparation  of    .     21 

Solubility- product 121 

quantitative      ....       7 
Solutions,  supersaturated     .     .     .122 

theories  of 115 

Standardization,  definition  of  .     .     64 

Standard  solutions 63,  64 

Stibnite,  analysis  of 108 

Stirring  rods 13 

Stoichiometry 132 

Strength  of  reagents 146 

Suciion,  use  of 14 

Sulphocyanite  process  for  silver  .  113 
Sulphur,  determination  of  .  .  32, 119 

Testing  of  washings 15 

Theories  of  solutions 115 

Theory  of  electrolytic  dissociation,  1 16 

Titration,  definition  of      ....  63 

Vacuum  pump,  use  of      ....     14 

Volumetric  analysis      .....     63 

definition  of   .       8 


Wash-bottles 9 

Washed  filters     .......     13 

Washing  of  precipitates  .     .     .  14,  122 

Washings,  testing  of 15 

Water,  ionization  of    .     .     .     124,  128 

Zinc,  determination  of      .     .     .     .     55 


